SCH3u Workbook
UNIT 1
MATTER, CHEMICAL TRENDS AND CHEMICAL BONDING
UNIT 1
MATTER AND CHEMICAL BONDING
1.1 The Study of Chemistry
Chemistry is the study of matter and its properties, the changes (chemical, physical, and nuclear) that matter experiences, and the energy that is associated with these changes.
The term STSE refers to science, technology, society, and the environment. These four areas are strongly interconnected. Science and technology, working on behalf of society, have produced chemicals that help every part of our lives. However, something that benefits society may harm the environment and end up harming society by endangering human health in some way. Science and technology working together have created a better quality of life for our society. At the same time, they have caused many environmental problems. Today, science and technology must work together to solve these and other problems.
1.2
Physical properties—colour, density, boiling point, melting point—are properties of matter that are observable without changing one type of matter into another type of matter. Chemical properties—reactivity (with oxygen, with acid, or with other substances), toxicity—are properties that are observable when one type of matter changes into a different kind of matter of different properties.
All measurements contain some degree of uncertainty because the last digit is estimated. The certainty of a calculated result depends on the certainty of the data, which in turn depends on the certainty (degree of accuracy) of the measuring instrument.
Calculated answers have the same number of significant digits as the least-certain measurement involved in the calculation; digits in excess of the number of least-certain digits must be rounded off when expressing the final answer. Precision refers to how close values are to each other, while accuracy refers to how close values are to an actual value.
Generally, the more precise a measurement is, the more significant digits it has.
Investigation 1-A Observing Aluminum Foil (text p. 13)
Complete Investigation 1-A and answer Analysis # 1, 2 Conclusion 3 a, b, c,
Using Measurement to Describe Matter
What is the name of the system of measurement that scientists rely on to communicate effectively. What does it allow scientists to do?
The system of measurement used by scientists is called the International System of Units or (S.I.). It allows scientists anywhere in the world to describe matter in the same quantitative language.
Name five of the base S.I. units?
Five of the base S.I. units are the metre (m), the kilogram (kg), the second (s), the mole (mol and the Kelvin (K).
Complete the following table on Important S.I. Quantities and Their Units Quantity Definition S.I. Units or their derived equivalents Equipment used to measure the quantity
mass
The amount of matter in an object kilogram (kg) gram (g) milligram (mg)
Balance or scale
length The distance between two points metre (m) centimeter (cm) millimetre (mm) ruler temperature The average kinetic energy of a substance
Kelvin )K) degrees Celsius(° C)
Thermometer
volume The amount of space an object occupies cubic metre (m3) cubic centimeter (cm3) litre (L) milliliter (mL)
beaker, graduated cylinder, pipette
Mole the amount of 6.023 x 1023 molecules of a substance mole (mol) calculated, not measured
Density the mass per unit volume of a substance kilograms per cubic metre (kg/m3) grams per cubic centimeter (g/cm3) calculated or measured
Energy the capacity to do work (to move matter)
Joule (J) calculated not measured
Measurement and Uncertainty
What are exact numbers?
Exact numbers are numbers you can count or are true by definition. They are numbers that you are certain about.
What two factors affect your ability to communicate measurements and calculations.
Two main factors affect your ability to communicate measurement and calculations. One factor is the instrument you use, and the other factor is your ability to read and interpret what the instrument is telling you.
Significant Digits, Certainty and Measurement
All measurements involve uncertainty. One source of this uncertainty is the measuring device itself. Another source is your ability to perceive and interpret a reading. The last digit (farthest right) in any measurement is always an estimate.
The digits that you record when you measure something are called significant digits. Significant digits include all the digits that you are certain about and a final, uncertain digit that you estimate.
How Can You Tell which Digits are Significant ?
Complete the following table of rules for determining significant digits:
Rules Examples
1. All non-zero numbers are significant 7.886 has four significant digit.
19.4 has three significant digits
527.266 992 \has nine significant digits
2. All zeros that are located between two non-zero numbers are significant 408 has three significant digits
25 047 has five significant digits
3. Zeros that are located to the left of a value are not significant 0.0907 has three significant digits. They are 9, the third 0 to the right, and the 7. The function of the 0.0 at the beginning is only to locate the decimal.
0.000 000 000 06 has only one significant digit
4. Zeros located to the right of a value may or may not be significant
22 700 may have three significant digits, or it may have five significant digits.
Explaining Three Significant Digits in the Example Above
The Great Lakes contains 22 700 km3 of water, but this an approximate value. Using scientific notation, you can rewrite 22 700 km3 as 2.27 x 104 km3 . This shows that only three digits are significant.
Explaining Five Significant Digits in the Example Above
If you were able to measure out the volume of water in the Great Lakes and verify the value of 22 700 km3 , then all five digits, including the zeros, would be significant. Once again, scientific notation lets you show clearly the five significant digits: 2.2700 x 104 km3
Complete Practice Problems #1, 2 (text p. 18)
Calculating With Significant Digits (text p. 20 – 22)
In this course, you will often take measurements and use them to calculate other quantities. You must be careful to keep track of which digits are significant.
Why is it important to keep track of significant digits?
It is important to keep track of significant digits because your results should not imply more certainty than your measure quantities justify. This is especially important when you use a calculator. Calculators usually report results with far more significant digits – than you data warrant. There are three rules for reporting significant digits in calculated answers.
Complete the rules for reporting significant digits in calculations in the table below:
Table 1 Rules for Reporting Significant Digits in Calculations
Rule 1 Multiplying and Dividing
The value with the fewest number of significant digits, going into the calculation, determines the number of significant digits that you should report in your answetr.
Rule 2 Adding and Subtracting
The value with the fewest number of decimal places going into the calculation determines the number of decimal places that you should report in your answer.
Rule 3 Rounding
To get the appropriate number of significant digits (rule 1)or decimal places (rule 2), you may need to round your answer.
If your answer ends in a number that is greater than 5, increase the preceding digit by 1. For example, 2.346 can be rounded to 2.35.
If your answer ends in a number that is less than 5, eave the preceding number unchanged. For example. 5.73 can be left as 5.7.
If your answer ends with 5, increase the preceding number by 1 if it is odd. Leave the preceding number unchanged if it is even. For example, 18.35 can be rounded to 18.4, but 18.25 is rounded to 18.2.
Sample Problem Reporting Volume Using Significant Digits
A student measured a regularly shaped sample o f iron and found it to be 6.78 cm long, 3.906 cm wide and 11 cm tall. Determine its volume to the correct number of significant digits.
Length = 6.78 cm (3 sig digits)
Width = 3.906 cm (4 sig digits)
Height = 11 cm (2 sig digits)
Solution
V = l x w x h = 6.78 cm x 3.906 cm x 11 cm
= 291.30948 cm3
The value 11 cm has the smallest number of significant digits, 2. Thus your answer can only have two significant digits. In order to have only two significant digits, you need to put your answer into scientific notation: v = 2.2 x 102 cm
Sample Problem Reporting Mass Using Significant Digits
Suppose that you measure the masses of four objects as 12.5 g, 145.67 g, 79.0 g, and 38.438 g.
What is the total mass of the four objects ?
Strategy:
• Add your masses together, aligning them at the decimal place.
• Underline the estimated (farthest right) digit in each value. This will help you keep track of the number of estimated digits in your final answer.
• In the question, two values have the fewest decimal places: 12.5 and 79.0.
• Round your answer so that it has only one decimal place.
12 .5 145.67 79.0 38.438 275.608 ∴ total mass of the objects is 275.6 g
(this makes sense because the answer has one decimal place, the same as the value in the question with the fewest decimal places
Complete Practice Problems # 3 a – g, p. 22 text)
Date: Name: Class:
Reinforcement Reporting Significant Digits in Calculations
Chapter 1
BLM 1-2
Goal
Reinforce your understanding of significant digits.
Procedure
1. State the number of significant figures in each of the following:
(a) 3570
(b) 17.505
(c) 41.400
(d) 0.51
(e) 0.000 572
(f) 0.009 00
(g) 41.50 × 10-4
(h) 0.007 160 × 105
(i) 1.234 00 × 108
(j) 0.000 410 0 × 107
2. Perform the following operations and give the answer to the correct number of significant digits.
(a) 15.1 + 75.32
(b) 178.904 56 – 125.805
(c) 4.55 × 10-5 – 3.1 x 10-5
(d) 0.000 159 + 4.0074
(e) 1.805 × 104 + 5.89 × 104
(f) 0.000 817 - 0.000 048 1
(g) 8.166 × 105 – 7.819 × 105
(h) 45.128 + 8.501 87 – 42.18
(i) 5.677 × 10-6 + 7.785 × 10-6
(j) 8.75 × 10-9 + 6.1157 × 10-9
(k) 1.99 ÷ 3.1
(l) 1200.0 ÷ 3.0
(m) 5.32 x 10-4 ÷ 4.218 × 10-8
(n) 45.32 x 2.3
(o) 0.024 00 ÷ 6.000
(p) 12.4 x 0.30
(q) (5.50 x 108) × (4 x 105)
(r) 7.4 ÷ 3
(s) 4.75 × 5
(t) 2.5 × 6.700 ÷ 0.891
To
Solutions to BLM1-2 Significant Digits
1.
(a) 3 or 4
(b) 3
(c) 5
(d) 2
(e) 3
(f) 3
(g) 4
(h) 4
(i) of 6
(j) 4
2. (a) 90.4
(b) 53.0991
(c) 1.5 ﾗ10 −5
(d) 4.0076
(e) 7.70 ﾗ10 4
(f) 0.000 770
(g) 0.347 ﾗ10 5
(h) 11.45
(i) 1.346 ﾗ10 −5
(j) 1.49 ﾗ10 −8
(k) 0.64
(l) 4.0 ﾗ10 2
(m) 1.26 ﾗ10 4
(n) 1.0 ﾗ10 2
(o) 4.000 ﾗ10 −3
(p) 3.7
(q) 2 ﾗ10 14
(r) 2
(s) 2 ﾗ10 1
(t) 19
CHAPTER 1 BLM ANSWER KEY
2.1 THE ATOMIC THEORY OF MATTER (Text p.p. 34 – 41) In 1809, John Dalton described atoms as solid, indestructible particles that make up all matter. Dalton’s model of the atom was often referred to as the _ - _ model.
Dalton’s concept of the atom was one of several ideas in his atomic theory of matter. Dalton’s atomic theory of matter (1809) has four main points:
• ________________________________________________________________________
________________________________________________________________________
• ________________________________________________________________________
________________________________________________________________________
• ________________________________________________________________________
________________________________________________________________________
• ________________________________________________________________________
________________________________________________________________________
________________________________________________________________________
• Law of Conservation of Mass: during a chemical reaction, the total mass of the substances involved does not change (mass of __________________________ equals the mass of the ______________________________
• Law of Definite Proportions: Elements always combine to form compounds in fixed _____________________________by mass. For example, pure water always contains the elements hydrogen and oxygen combine in the following proportions: 11 % hydrogen, and 89% oxygen
Chemistry Background
What is a physical property ?
Physical properties—colour, density, boiling point, melting point—are properties of matter that are observable without changing one type of matter into another type of matter. What is a chemical property? Chemical properties—reactivity (with oxygen, with acid, or with other substances), toxicity—are properties that are observable when one type of matter changes into a different kind of matter of different properties.
THE MODERN VIEW OF THE ATOM An is the smallest particle of an element that still retains the identity and properties of the element. For example, the smallest particle of the writing material in a pencil is a carbon atom. Pencil “lead” is actually a substance called graphite. Graphite is a form of the element carbon. An average atom is about 10 – 10m in diameter. Atoms themselves are made up of even smaller particles. These subatomic particles are , , and . Protons and neutrons cluster together to form the central core, or , of an atom. Fast moving electrons occupy the space that surrounds the nucleus of the atom. As their name implies, subatomic particles are associated with electrical charges as shown in the table below:
Subatomic Particle
Charge
Symbol
Mass (g)
Radius (m)
Electron
Proton
Neutron 0
N0
1.67 x 10- 24
10 – 15
EXPRESSING THE MASS OF SUBATOMIC PARTICLES Subatomic particles are incredibly small. For example one proton or a neutron has a mass of g. It is inconvenient to measure mass of subatomic particles using units such as grams. Instead, chemists use a unit called an atomic mass unit (symbol ___). A proton has a mass of about 1 u, that is equal to 1.66 x 10 – 24 g. A proton is about 1837 times more massive than an electron. From the table above, the mass of an electron is 9.02 x 10 – 28. This value is so small that scientists consider the mass of an electron to be approximately equal to zero. Thus, electrons are not taken into account when calculating the mass of an atom.
EXPRESSING THE MASS IF SUBATOMIC SUBSTANCES
• Atomic Mass Unit
The Nucleus of an Atom
• Atomic Number (_______)
• Mass Number (______)
Information about an element’s protons and neutrons is often summarized using the chemical notation shown in Figure 2.3. Mass # A
X
Atomic # Z
In the figure above, which letter represents the atomic symbol of the element ? (X)
In the following example, identify the mass number and the atomic number
19 F 9
What do these numbers indicate ?
The mass umber (the superscript 19) indicates that fluorine has a total of 19 protons and neutrons. The atomic number (the subscript 9) indicates that fluorine has 9 protons.
How can you determine the number of protons in the nucleus of an atom?
To find the number of neutrons, subtract the atomic number from the mass number.
Number of Neutrons = mass # - atomic #
Or
Number of Neutrons = A - Z
How many neutrons would this isotope of fluorine have?
Number of Neutrons = A - Z = 19 - 9 = 10
PRACTICE PROBLEMS
1. Complete the missing values or chemical notations about the atoms from the table on p. 37 of your textbook below.
in a. b. c.
d. e f.
g h i
j k l
m n o
p. q. r.
s. t. u.
v. w. x.
When chemists refer symbolically to an element, they will often leave out the atomic number. For example, for oxygen, they would write 16 O or O – 16.
USING THE ATOMIC NUMBER TO INFER THE NUMBER OF ELECTRONS
The atomic number and mass number do not give you direct information about the number of neutrons in an element. They do not give you the number of electrons either. You can infer the number of electrons, however, from the atomic number. The atoms of each element are electrically neutral. This means that the positive charges (protons) and negative charges (electrons) must balance one another. Therefore, in a neutral atom of any element, the number of protons is equal to the number of electrons. For example, a neutral hydrogen atom contains one proton, so it must contain one electron. A neutral oxygen atom contains six protons, so it must contain six electrons.
ISOTOPES AND ATOMIC MASS All neutral atoms of the same element contain the same number of protons and, therefore, the same number of electrons. The number of neutrons can vary, however. For example, most of the oxygen atoms in nature have eight neutrons in their atomic nuclei. In other words, most oxygen atoms have a mass number of 16 (8 protons + 8 neutrons). As you can see in the figure below there are also two other naturally occurring forms of oxygen. One of these has nine neutrons, so A = 17, and the other has ten neutrons, so A = 18. These different forms of oxygen are called . Isotopes are ______
Fig. 2.4
p. 38 Text
The isotopes of an element have very chemical properties because they have the same number of protons and electrons. They differ in , however, because they have a different number of neutrons. Some isotopes are more unstable than others. Their nuclei are more likely to decay, releasing energy and subatomic particles. This process, called ______, happens spontaneously. All uranium isotopes, for example, have unstable nuclei. They are called radioisotopes. May isotopes are not radioisotopes. Oxygen’s three naturally occurring isotopes, for example, are stable. In contrast, chemists have successfully synthesized ten other isotopes of oxygen, all of which are unstable radioisotopes. Radioisotopes decay because their nuclei are unstable. The time it takes for a nucleus to decay varies greatly. For example, it takes billions of years for only half of the nucleus of naturally occurring uranium – 238 to decay. The nuclei of other radioisotopes, mainly those synthesized by scientists, decay much more rapidly. The nuclei of some isotopes, such as sodium – 22, take about 20 years to decay. For calcium – 47, this decay occurs in a mater of days. The nuclei of most radioisotopes decay so quickly, however, that the radioisotopes exist for mere fractions of a second. Atomic Weights
This is also called the “Average Atomic Mass”. It is the weighted average of all naturally occurring isotopes of an element. A.W. = (% 1) (m 1) + (% 2) (m 2) + (% 3) (m 3) . . . . . . . . 100
where % = the percent occurrence of the isotope in nature where m = the mass of the isotope
Example 1 Sulfur has three isotopes which have the following relative abundances: S – 32 = 95% S – 33 = 0.8 % S – 34 = 4.2 % Calculate the atomic weight of sulfur.
ELECTRONS IN ATOMS
What is the importance of the electron to the atom ? Recall that electrons occupy the space surrounding the nucleus therefore they are the first subatomic particles that are likely to interact when atoms come near one another. It is the and ________ of electrons that determine how an atom will react, if at all, with other atoms. ______are responsible for the chemical properties of elements.
REVISITNG THE ATOMIC THEORY John Dalton did not even know about subatomic particles when he developed his atomic theory. Even so, the modern atomic theory, shown below, retains many of Dalton’s ideas, with only a few modifications. Not all chemists believed that Dalton’s atoms existed. Other scientists considered atoms to be a valuable idea for understanding matter and its behavior. They did not , however, believe that atoms had any physical reality. The discovery of electrons (and later, the other subatomic particles) finally convinced scientists that atoms were more than simply an idea. Atoms, they realized, must be ______. The atomic theory is a landmark achievement in the history of chemistry. It shaped the way that all scientists, especially chemists, think about matter.
Table The modern Atomic Theory
p. 39 Text
SECTION REVIEW
1. For the table below, fill in the missing information in the appropriate space provided.
ELEMENT ATOMIC NUNMBER
MASS NUMBER
NUMBER OF PROTONS
NUMBER
OF ELECTRONS
NUMBER OF NEUTRONS
(a)
(b)
108
(c)
47
(d)
(e)
(f)
(g)
33
(h)
42
(i)
35
(j)
(k)
(l)
45
(m)
79
179
(n)
(o)
(p)
(q)
(r)
(s)
(t)
50
69
1. a. b. c. d.
e f. g. h.
i. j. k. l.
m. n. o. p.
q. r. s. t.
2. Explain the difference between an isotope and a radioisotope. Provide an example other than oxygen to support your answer.
3. Examine the information presented in the following pairs
3 3 14 16 19 18 H and He C and N F and F 1 2 6 7 9 9
(a) For each pair, do both members have the same number of protons ? electrons ? neutrons?
(b) Which pair or pairs consist of atoms that have the same value for Z? Which consists of atoms that have the same value for A ?
4. Compare Dalton’s atomic theory with the modern atomic theory. Explain why scientists modified Dalton’s theory.
5. Fill in the following table:
Protons
Electrons
Neutrons
P – 31
Al – 27
H-1 (protium)
H – 2 (deuterium)
NOTES
NOTES
6.
# 15, & 17 FROM Sch3a workbook 2000
Unit 2
7.
EXPANDING THE MODEL OF THE ATOM (p.p. 656 – 659)
Emission Spectra
Each wavelength of visible light is associated with a ______. When white light is shone through a gas discharge tube, it produces a bright line ______. A bright line spectrum is a series of narrow lines having specific colour ( _____), separated by colourless space. Each element produces a different and characteristic colour spectrum. You can see several examples of these spectra in the figure below. Scientists who were studying the atom, needed to take emission spectra into account.
Fig. D1
p. 656 text
Bohr’s Model of the Atom In 1913, Danish physicist, Niels Bohr developed a model of the atom that explained hydrogen emission spectrum. In Bohr’s model electrons ______ the nucleus in the same way that the Earth orbits the sun, as shown in the figure below.
Fig. D-2
p. 656 Text
The following three points of Bohr’s theory help explain hydrogen’s emission spectrum:
• In Bohr’s model, atoms have specific allowable ______levels. He called these energy levels stationary states. Each of these levels corresponds to a fixed, circular orbit around the nucleus.
• An atom does not give off energy while its electrons are in a stationary state
• An atom changes states by giving off, or absorbing a quantity of light energy exactly equal to the ______in energy between the two stationary states.
Bohr’s model was revolutionary because he proposed that the energy absorbed or emitted by an atom needed to have ______ ______. The energy change was ______, rather than continuous. When something is quantized, it is limited to discrete amounts or multiples of discrete amounts. Two great scientists paved the way for this surprising idea. German physicist Max Plank had already proposed that light was quantized, meaning that it exists in packets. Building on this idea, Einstein proposed that light could behave as particles, which he called photons. The energy associated with the light in a line spectrum corresponds with the change in energy of an electron as it moves _______or _______________ an energy level. For example, when electrons in hydrogen atoms that have been excited to the third energy level and subsequently drop to the second energy level, they emit light that has a specific wavelength. They emit photons of _____________ light that have a wavelength of 656.5 nm. These photons cause the red line spectrum for hydrogen seen in Figure D1 above. Why A New Model ? Bohr’s model worked fine at predicitng the line spectrum for hydrogen or for ions with only one electron e.g. He +, Li 2+ , Be 3+ etc. The model failed, however, when it was applied to the emission spectra of atoms that had more than one electron. Bohr’s model needed to be modified because it was too simple to explain the experimental evidence. Sublevels The spectra of many electron atoms suggested that a more complex structure was needed. Notice that in Figure D1, that spectra for these more complex atoms have groups of lines close together. The groups of lines are separated by spaces. The large spaces represent energy differences between ______ ______, while the smaller spaces represent energy ______within levels. If the electrons are changing energy within the levels, this suggests that there are sublevels within each level., each with its own slightly different energy. The idea of different energy levels 1, 2, 3, 4, . . . remains, but each energy level is split up into sublevels called _____, _____, _____, _____. Examine the table below to see how electrons are arranged in sublevels for energy levels 1 to 4.
Table D1
p. 657 text Each sublevel has its own energy. Examine the figure on the following page to see the relative energies of the different levels and sublevels. But where would you find electrons in this new model ? How are the sublevels oriented in space ?
Particles with Wave Like Properties
In 1924, Louis de Broglie suggested that all matter had -like properties. De Broglie developed an equation that allowed him to calculate the wavelength associated with any object. Objects we can see and interact with have calculated wavelengths that are smaller than electrons. Their wavelengths are so tiny compared to their size, that they do not have any measurable effect on the motion of the objects
For very tiny moving particles such as electrons, however, the wavelength becomes very significant. De Broglie’s theory was proven by experiment when streams of electrons produced diffraction patterns similar to electromagnetic radiation that was already known to travel in waves.
Orbitals
In 1920, Erwin Schrodinger used de Broglie’s idea that matter has wavelike properties. Schrodinger proposed the quantum mechanical model of the atom. In this new model, he abandoned the notion of the electron as a small particle orbiting the nucleus. Instead, he described the behaviour of electrons in terms of wave functions.
This imprecise nature of Schrodinger’s model was supported shortly afterwards by principle proposed by Heisenburg. Heisneburg demonstrated that it is impossible to know both an electron’s path and its exact location. Therefore it is impossible to know the exact location of an electron at any time, without some degree of uncertainty. Heisenburg’s uncertainty principle shows that you can never know both the ______and the ______ of an electron beyond a certain measure of precision. Heisneburg also showed that if you could know either the velocity or position precisely, then the other property would be uncertain. Therefore, when talking about where electrons are found, you can only talk in terms of
______ ______.
Schrodinger used a mathematical wave equation to define the probability of finding an electron within an atom. There are multiple solutions to this wave equation, and Schrodinger called these solutions ______. Each solution provides information about the energy and location of an electron within an atom. Each orbital has a specific energy associated with it, and each contains information about where inside the atom, the electrons would spend most of their time.
These regions of space where an electron is most likely to be found are called ______. In solving the wave equation, three unknowns were encountered (a fourth was added later). These unknowns are called the ______ ______and describe different properties of the electron.
The Four Quantum Numbers i. The Principle Quantum Number (_____)
• Indicates the ______an electron has
• Tells us which energy ______the electron is in
• Can have values of n = ______
• The higher the ‘n’ value, the ______the energy
• Correspond to Bohr’s orbits
2. The ______Quantum Number (_____)
• Indicates the ______of the orbital
• It can have values of L = ______ ______
• n- value possible ‘L’ values
1
2
3
4.
Note: the total number of possible ‘L’ values is the same as the ‘n’ value L- VALUE
0
1
2
3
4
SHAPE
SPEHRICAL
DUMBBELL
DOUBLE DUMBBELL
??
??
NAME
____– orbital
____– orbital
____– orbital
____– orbital
____– orbital
3. The _____________________Quantum Number (_____)
• indicates the _____________________of the orbitals about the nucleus.
• The number of values indicates the number of orbitals that exist with a given
• M L can have the values of ____________________________________ L- VALUE
0
1
2
3
4
SHAPE
SPEHRICAL
DUMBBELL
DOUBLE DUMBBELL
??
??
NAME
S – orbital
P – orbital
D – orbital
F – oribital
G - orbital
3. The Magnetic Quantum Number (M L)
• indicates the orientation of the orbitals about the nucleus.
• The number of values indicates the number of orbitals that exist with a given
• M L can have the values of - L . . . 0 . . . + L
L
0
1
2
3
4
5
NAME
S
P
D
F
G
H
M L
0
-1, 0, 1
-2, -1, 0, 1, 2
-3, -2, -1, 0, 1, 2, 3
NUMBER OF POSSIBLE ORBITALS
1
3
5
7
4. _________________Quantum Number (_____)
• Related to the _________________ of an electron
• All electrons spin on their axis. The two different spins are assigned the value of
________________________
• indicates that each orbital can hold a maximum of ________ electrons. If there are two electrons in an orbital, they must have opposite spins.
n
L
NAME
M L # of
Orbitals
M S # of
Electrons Total # of
Electrons
1
2
2
3
3
3
4
4
4
4
The ‘n’ Rules
• The nimber of different shapes (or ) of orbitals in a given energy level =
• The total number of orbitals in a given energy level =
• The total number of electrons in a given energy level =
Shapes of Orbitals The shapes of the probability graphs of Schrodinger’s wave functions are the shapes of the orbitals in which electrons reside in atoms. You can visualize orbitals as electron clouds. The shape of each cloud is based on probability – it tells where the electron spends most of its time. Examine the figure below. The ‘s’ orbitals are spherical in shape. Each of these spherical shells contains two electrons. There are three ‘p’ orbitals for each sublevel, each with a capacity for two electrons. Each ‘p’ orbital is shaped something like a dumbbell. For each subleve, the ‘p’ orbitals are oriented along the x, y, and z axes. The d’ and ‘f’ orbitals are quite complex in shape. Each ‘d’ sublevel contains 5 ‘d’ orbitals and each ‘f’ sublevel contains 7 ‘f’ orbitals.
S & p orbitals
SCH3a 2000 workbook
Quantum unit
D orbital diagram
Sch3a workbook 2000
f orbital diagram
Sch3a workbook 2000
Filling Orbitals
Each orbital can hold a maximum of ________ electrons. A hypothesis to suggest why this is true suggests that electrons spin around their own axes as they move around the nucleus, generating magnetic fields. They can spin in a positive direction or a negative direction . Electrons with opposite spins attract each other. This attraction partially counteracts the repulsion between two negatively charge electrons. In 1925, Pauli proposed that only two electrons of opposite spin could occupy an oribital.
The Pauli Exclusion Formula - states that ________________________________________________________________________
How do electrons fill orbitals within atoms? They do so to minimize the potential energy of the atom.
1. The maximum number of electrons that any orbital can hold is _____________
2. They will fill the ________________________energy level first. The 1s orbital will fill before the 2s orbital, which fills before the 2 p orbitals and so on.
1. When occupying two or more orbitals with the same energy (for example any of the 3 ‘p’ orbitals), electrons will half fill each orbital until all are half filled before adding a second electron to each one. This is called ________________________Rule. You can show how electrons fill orbitals using superscripts in a notation called the electron configuration of the atom. For example, a boron atom contains 5 electrons. The electron configuration for Boron would be 1s 2 2s 2 2p 1. The electron configuration for Nitrogen is 1s 2 2s 2 2p 3.
First half of Periodic Table & Their Quantum numbers
SCH3A Workbook 2000
Electron Configurations
2nd half of Periodic Table & Their Quantum numbers
SCH3A Workbook 2000
1. 21 Sc
2. 23 V
3. 30 Zn
4. 35 Br
5. 50 Sn
6. 57 La
7. 56 Ce
8. 64 Gd
9. 50 Hg
10. 84 Po
A Schematic Energy Level diagram of a many electron atom
Whole page
Sch3a 2000
Quantum Unit
Whole page
Quantum Review 1
Quantum Review 1 page 2
Notes
Notes
ATOMS, ELEMENTS, AND THE PERIODIC TABLE (Text p.p. 40 – 47) By the mid 1800’s, there were 65 known elements. Chemists studied these elements intensively and recorded detailed information about their reactivity and the masses of their atoms. Some chemists began to recognize patterns and behaviour of many of these elements.
Fig 2.5
p. 40
Other sets of elements display similar trends in their properties and behaviour. For example, oxygen (O), sulfur (S), selenium (Se), and tellurium (Te) share similar properties. The same is true for fluorine (F), chlorine (Cl), bromine (Br), and iodine (I). These similarities prompted chemists to search for a fundamental property that could be used to organize all elements. One chemist, Dmitri Mendeleev (1834 – 19907) sequenced the known elements in order of increasing atomic mass. The result was a table of the elements, organized so that elements with similar properties were arranged in the same ________________________. Because Mendeleev’s arrangement highlighted the periodic (repeating) patterns of properties, it was called a periodic table. The term ________________________means “repeating in an identifiable pattern.
The modern periodic table is a modification of the arrangement first proposed by Mendeleev. Instead, of organizing the elements in order of increasing atomic _____________, the modern periodic table organizes elements according to atomic ________________________. According to the Periodic Law, the chemical and physical properties of the elements repeat in a regular, periodic pattern when they are arranged according to their atomic number.
p. 41 text PRACTICE PROBLEMS
2. Identify the name and symbol of the elements in the following location of the periodic table:
a. Group 14 (IVA), Period 2
b. Group 11 (IB), Period 4
c. Group 18 (VIIIA), Period 6
d. Group 1 (IA), Period 3
e. Group 12 (IIB), Period 5
f. Group 2 (IIA), Period 4
g. Group 17, (VIIA), Period 5
h. Group 13 (IIA), Period 3
_____________________________
_____________________________
_____________________________
_____________________________
_____________________________
_____________________________
_____________________________
_____________________________
ELECTRONS AND THE PERIODIC TABLE
We have seen how the periodic table organizes elements so that those with similar properties are in the same group. You have also seen how the periodic table shows a clear distinction among metals, non-metals, and metalloids. Other details of the organization of the periodic table may seem strange. Why, for example, are there different numbers of elements in the periods.
The reason for this, and other details of the periodic table’s organization, involves the number and arrangement of ________________________in atoms of each element.
ELECTRONS AND ENERGY LEVELS Electrons cannot move haphazardly. Their movement around an atomic nucleus is restricted to fixed regions of space. These regions of space are three dimensional. The following figure shows a representation of these regions.
Fig 2.8
p. 43 text
These regions are not solid, but rather are volumes of space in which electrons may be found. They are referred to as energy shells or energy levels. An electron that is moving in a lower energy level is found ________________________to the nucleus. It would have less energy than it would if it were moving in a higher energy level. There is a limit to the number of electrons that can occupy each energy level. For example, a maximum of _______________ electrons can occupy the first energy level. A maximum of ________________________electrons can occupy the second energy level. The periodic trends that result from organizing the elements by their atomic number are linked to the way in which electrons occupy and fill energy levels.
Text
Fig 2.9
p. 44
As shown in the figure above, a common way to show the arrangement of electron in atoms is to draw circles around the atomic symbol. Each circle represents an energy level. Dots represent electrons that occupy each energy level. This kind of diagram is called a Bohr –Rutherford diagram. Figure 2.9 B shows that the first energy level is full when _____________ electrons occupy it. Only two elements have two or fewer electrons: hydrogen and helium. Hydrogen has one electron, and helium has two. These elements, with their electrons in the first energy level, make up Period _______ of the periodic table. As you can see in Fig. 2.9 C, Period 2 elements have two energy levels. The second energy level is full when ________electrons occupy it. Neon, with a total of ten electrons, has its first and second energy levels filled. Notice how the second energy level fills with electrons as you move across the period from lithium to fluorine.
PATTERNS BASED ON ENERGY LEVELS AND ELECTRON ARRANGEMENTS The structure of the periodic table is closely related to ________________________ levels and the arrangement of electrons. Two important patterns result from this relationship. One involves periods, and the other involves groups. The Period Related Pattern
As you can see in Figure 2.9, elements in Period 1 have electrons in one energy level. Elements in Period 2 have electrons in two energy levels. This pattern applies to all seven periods. An element’s period number is the same as the number of ________________________levels that the electrons in the atom occupy. Thus, you could predict that Period 5 elements occupy 5 different energy levels. What about the inner transition elements – the elements that are below the periodic table? Figure 2.10 shows how this pattern applies to them. Elements 58 through 71 belong in Period 6, so their electrons occupy six different energy levels. Elements 90 through 103 belong in Period 7, so their electrons occupy seven different energy levels. Chemists and chemical technologists tend to use only a few of the inner transition elements on a regular basis. Thus it is more convenient to place all inner transition elements below the periodic table.
The Group Related Pattern The second pattern emerges when you consider the electron arrangements in the main- group elements: the elements in Group 1 (1A), 2 (2A) and 13 (3A) to 18 (8A). All the elements in each main group have the same number of electrons in their highest (outer) energy level. The electrons that occupy the outermost energy level are called ________________________ electrons. It is these valence electrons that are involved when atoms form compounds. In other words, ________________________ ________________________are responsible for the chemical behaviour of elements. You can infer the number of valence electrons in any main group element from its group number. For example, Group 1 (1A) elements have _____ valence electron. Group 2 (2A) elements have _____ valence electrons. For elements in Group 13 (3A) to 18 (18A), the number of valence electrons is the same as the second digit in the current numbering system. It is the same as the number only in the old numbering system (in brackets). For example, Group 15 (5A) have five valence electrons while Group 17 (7A) have seven valence electrons.
USING LEWIS STRUCTURES TO REPRESENT VALENCE ELECTRONS It is time consuming to draw electron arrangements using Bohr-Rughterford diagrams. It is much simpler to draw Lewis structures to represent elements and their valence electrons. To draw a Lewis structure, you replace the nucleus and the inner energy levels of an atom with its atomic symbol. Then you place dots around the atomic symbol to represent the valence electrons. The order you place the first four dots is up to you, however, place a dot on each side of the symbol before doubling up on any dot.
Fig 2.11
Text
Practice Problems
2. On the blank periodic table, sketch the electron arrangements for the first 20 elements using Lewis structures.
3. Use the periodic table to draw Lewis structures for the following elements: barium (Ba), gallium (Ga), tin (Sn), bismuth (B), iodine (I), cesium (Cs), krypton (Kr) and xenon (Xe)
4. Identify the number of valence electrons in the outer energy levels of the following elements:
a. chlorine f. lead
b. helium g. antimony
c. indium h. selenium
d. strontium i. arsenic
e. rubidium j. xenon NOTES
THE SIGNIFICANCE OF A FULL OUTER ENERGY LEVEL The ________________________ ________________________in Group 18 (VIIIA) are the only elements that exist as individual atoms in nature. They are extremely ________________________. They do not naturally form compounds with other atoms. What is it about noble gases that explains their behaviour ? Recall that chemical reactivity is determined by ________________________electrons. Thus, there must be something about the arrangement of the electrons in noble gases that explains how unreactive they are. All the noble gases have outer energy levels that are completely ________________________with the maximum number of electrons. Helium has a full outer energy level of ______ valence electrons. The other noble gasses have ______ valence electrons in the outer energy level. Chemists reason that having a full outer energy level must be a very ________________________electron arrangement. What does this stability mean ? It means that a full outer energy level is unlikely to change. Scientists have observed that, in nature, situations or systems of lower energy are favored over situations of higher energy. For example, a book on a high shelf has more potential energy than a book on a lower shelf. If you move the book to the ground, it has low potential energy, that is, it is more stable. When atoms have eight electrons in their outermost energy level, (or two for hydrogen and helium), chemists say they have a stable ________________________. An octet is a very stable electron arrangement. Summary
You have seen that the structure of the periodic table is directly related to energy levels and arrangement of electrons. The patterns that emerge from this relationship enable you to predict the number of valence electrons for any main group element. They also enable you to predict the number of energy levels that an element’s electrons occupy.
Orbital Box Notation Using this method, each orbital is shown as a box with ________________________ representing the electrons. If there are two electrons in the same orbital box, then they must have opposite spins. This is shown by having the arrows facing in different _____________________.
Element Electron Configuration Orbital Box
4 Be
6 C
7 N
Hund’s Rule _______________________________________________________________________________________________________________________
8 O
9 F
21 Sc
50 Sn
Ground State Electron Configurations with this arrangement, the electrons are in the ________________________possible orbitals. This is the order given by the Periodic Table.
Excited State Configurations In order to obtain these configurations, energy must be added to the atom. When energy is added, one or more electrons is moved to a higher energy orbital. To which orbital the electron gets moved depends on the amount of energy that is added.
12 Mg Ground State Electron Configuration Excited State (one of many)
Electron Configurations and Valences Atoms react with other atoms to become more ________________________ (to have a more stable electron arrangement). It is the arrangement of the atom’s outermost electron that determine its chemical properties. To become more stable, an atom can either _______________ or ___________________ electrons. If an atom is giving away electrons, then the ones with the highest ______ value always go first. Note that the rule for taking electrons away does not follow the rule for adding electrons to an atom. Stable Arrangements There are several different arrangements which can make an atom more stable:
1. The most stable arrangement is achieved by having a completely filled energy level.
Examples:
2. The second most stable configuration involves having fully occupied o completely empty energy levels. A sub-energy level refers to all the orbitals of one-kind at that energy level.
Examples:
3. The third most stable configuration is achieved by having half-filled sub-energy levels Examples:
Valences Predict the most likely valences for the following:
1. 7 N
2. 16 S
3. 30 Zn
4. 50 Sn
Explain why:
i. Fe (# 26) has a (+2) and a (+3) valence. Which is more stable ?
ii. Rh (# 45) has a (+2) and a (+4) valence.
iii. Cerium (# 58) has a (+3 and a (+4) valence
Exceptions to the Predicted Electronic Configurations There are two families whose electron configurations are different than what would be predicted. They are the families VIB and IB. The elements in these families spontaneously rearrange their electrons to make themselves more stable. To do this, they promote an electron from an ‘s’ orbital and excite it to a ‘d’ orbital.
Examples
24 Cr
42 Mo
29 Cu
47 Ag
Chemical Rections and the desire to become more stable
Sch 3a Workbook 2000
Section Review
1. State the Periodic Law and provide at least two examples to illustrate its meaning.
2. Identify the Group number for each of these sets of elements. Then write the symbols for the elements within it:
• alkali metals
• noble gases
• halogens
• alkaline earth metals
3. Identify the elements that is described by the following information. Refer to a periodic table as necessary.
• It is a Group 14 (IIIA) metalloid in the third period
• It is a Group 15 (VA) metalloid in the fifth period
• It is the other metalloid in Group 15
• It is a halogen that exists in the liquid state at room temperature
4. What is the relationship between electron arrangement and the organization of the elements of the periodic table ?
5 How many valence electrons are there in an atom of each of these elements ?
(a) neon sodium magnesium bromine chlorine silicon sulfur helium strontium tin
(b) Present your answers from Part A in the form of Lewis structures.
(c) Classify each element from Part A as a metal, non-metal, or a metalloid
6. Compare and contrast noble gases with other elements.
7.
(a) Draw Lewis Structures for each of the following elements: lithium potassium magnesium aluminum carbon
(b) Which of these elements have the same number of energy levels ?
(c) Which have the same number of valence electrons
8. Identify the elements with the following electron distributions by orbits:
a. 2, 8, 8 b. 2, 8, 6 c. 2, 4 d. 2, 8, 8, 2
5. What is meant by a stable octet ?
6. How many atoms must each of the following elements gain to achieve a stable octet ?
a. fluorine b. oxygen c. nitrogen d. carbon
7. How many electrons must the following atoms lose in order to achieve a stable octet
a. silicon b. phosphorus c. sodium d. argon
8. How does the sign of the valence relate to the gain or loss of electrons ?
9. Explain why atoms form ions.
Notes
Notes
Notes
Notes
Notes
2.3
PERIODIC TRENDS INVOLVING THE SIZES AND ENERGY LEVELS OF ATOMS In Section 2.1, we learned that the size of a typical atom is about 10 – 10 m, however, we also know that the atoms of each element are distinctly different. For example, the atoms of different elements have different numbers of protons. This means of course, that they also have different numbers of electrons. You may predict that the size of an atom is related to the number of protons and electrons it has. Is there any evidence to support this prediction, If so, is there a pattern that could help you to predict the relative size of an atom for any element on the periodic table ? Chemists define, and measure, an atom’s size in terms of its radius. The radius of an atom is the distance from its nucleus to the approximate outer boundary of the cloud-like region of its electrons. This boundary is approximate because atoms are not solid spheres. They do not have an outer boundary.
Fig. 2.12
p. 49 Text
Inv 2-A
p.p. 50 & 51
Text
Inv 2-A
p.p. 50 & 51
Text
Trends for Atomic Size (Radius) There are two general trends for atomic size:
• As you go down each group in the periodic table, the size of an atom ___________________. This makes sense if you consider energy levels. As you go down a group, the valence electrons occupy an energy level that is further from the nucleus. Thus, the valence electrons experience less attraction for the nucleus. In addition, electrons in the inner energy levels block, or shield the valence electrons from the attraction of the nucleus. As a result, the total volume of the atom, and thus size, increase with each additional energy level.
• As you go across a period, the size of an atom ________________________. This trend might seen surprising at first, since the number of electrons increases as you go across a period. The size of the atom decreases, however, because the positive charge of the nucleus also increases across a period. As a result, the positive force exerted by the nucleus increases, pulling the outer electrons ________________________, thus reducing the atom’s total size
Fig 2.13
p. 552 Text
General Trends for Atomic Radius Graph
SCH 3A workbook 2000
Trends on he P.T.
Sch3a Workbook
Practice Problems
7. Using only their location in the periodic table, rank the atoms in each set by decreasing atomic size. Explain your rankings. (a) Mg Be Ba
(b) Ca Se Ga
(c) Br Rb Kr
(d) Se Br Ca
(e) Ba Sr Cs
(f) Se Br Cl
(g) Mg Ca Li
(h) Sr Te Se
(i) In Br I
(j) S Se O Notes
Trends for Ionization Energy A neutral atom contains an equal number of positive charges (protons), and negative charges (electrons). It takes ________________________to overcome the attractive force of the nucleus and pull an electron away from a neutral atom. The energy required to remove an electron from an atom is called ________________________energy. The particle that results when a neutral atom gains electrons or gives up electrons is called an ion. Thus, an ion is a ________________________particle. An ion that ________________________electrons becomes a negatively charged anion, while an ion that ________________________up electrons becomes a positively charged cation. The figure below shows the formation of ions for several elements. As you examine the diagrams, pay close attention to
• the energy level from which the electrons are gained or given up
• the charge of the ion that is formed when an atom gains or gives up electrons
• the arrangement of the electrons that remain after the electrons are gained or given up.
Fig . 2.14
p. 53 text
If you try to visualize the periodic table as a cylinder, (see Fig. 2.15 below) rather than a flat plane, you can begin to see the relationship between ion formation and the electron arrangement of the noble gases.
Fig 2.15
p. 54 Text
The metals that are the main-group elements tend to ______________ up electrons and form ions that have the same number of electrons as the nearest ________________________gas. Non-metals tend to _____________________ electrons and form ions that have the same number of electrons as the nearest noble gas. For example, when a sodium atom gives up its single valence electron, it becomes a ________________________charged sodium ion. Its outer electron arrangement is like ________________________outer electron arrangement. When a fluorine atom gains an electron, it becomes a ________________________charged ion with an outer electron arrangement like that of ________________________. Figure 2.15 can help you determine the charge on an ion. Count the number of groups an ion is from the nearest noble gas. That number is the ________________________on the ion. For example, aluminum is three groups away from neon, and so aluminum has a charge of + 3. Sulfur is two groups away from argon, thus sulfur has a charge of 2 -. Remember metals form positive ions (________________) while non-metals can form both positive ions (________________) or negative ions (________________). Negative ions are formed by non-metals when they combine with metals. One of the reasons that the periodic table is useful is that it can help us remember the types of ions formed by many of the representative elements. For example, except for hydrogen, the neutral atoms of Group 1A always lose one electron each when they react, thereby becoming ions with a charge of 1 +. Similarly, elements in Group II always lose two electrons to become ions with a charge of 2+. In Group III, the only important ion that we need to consider at this time is aluminum, which loses three electrons to form an ion with a charge of 3 +. Notice that the charges on the cations is the same as their traditional Group numbers on the periodic table. Although this generalization does not work for all metallic elements (it doesn’t work for the transition elements), it does help us remember what happens to the members of Group I, Group II, and aluminum when they react, Among the non-metals on the right side of the periodic table, we also find some useful generalizations. For example, when they combine with metals, the_____________________ (Group VIIA or Group 17) form ions with a 1 – charge, the non-metals in Group VIA or Group 16, form ions with a 2 – charge. Notice that the number of negative charges on the anion is equal to the number of spaces to the right that we have to move on the periodic table to get to a noble gas.
Littlediagram from new sjuce
Middle of page
The figure below shows the ionization energy that is required to remove one electron from the outer energy level of atoms of the main group elements This energy is called the ____________-ionization energy. It is measured in units of kJ / mol. As you can see, atoms that give up electrons easily have _______________ ionization energies. You would probably predict that the alkali metals of Group 1 (1A) would have low ionization energies. These elements are, in fact, extremely reactive because it takes so little energy to remove their single valence electron.
Fig. 2.16
p. 54 text
All elements, except _________________________, have more than one electron that can be removed. Therefore, they have more than one ionization energy. The energy that is need to remove a second electron is called the ________________________ionization energy. The energy needed to remove a third electron is called the ________________________ionization energy, and so on. Summarizing Trends for Ionization Energy
Upon examination of the periodic trend for ionization energy it is the opposite of the trend for atomic radius.
Fig. 2.17
Text p. 55
Bar Graph from SCH3a workbook 2000
Atomic Radius Vs I.E.
Although there are a few exceptions, there are two general trends for ionization energy:
• Ionization energy tends to ________________________down a group as electrons in the outer energy level are further from the positive force of the nucleus, thus they are easier to remove than electrons in lower energy levels.
• Ionization energy tends to ________________________across a period. As you go across a period, the attraction between the nucleus and the electrons in the outer energy level increases. Thus, ______________________ energy is required to pull an electron away from its atom. For this trend to be true, you would expect the noble gases to have the highest ionization energies, and as is shown in Figure 2.16, they do. Practice Problems
8. Using only a periodic table, rank the elements in each set by increasing ionization enegy. Explain your answers. Explanation
(a) Xe He Ar
(b) Sn In Sb
(c) Sr Ca Ba
(d) Kr Br K
(e) K Ca Rb
(f) Kr Br Rb
9. Using only a periodic table, identify the atom in each of the following pairs with the lower first ionization energy. (a) B O
(b) B In
(c) I F
(d) F N
(e) Ca K
(f) B Tl Trends for Electron Affinity Electron affinity is the change in ________________________that occurs when an electron is added to the outer energy level of an atom to form a negative ion. Figure 2.18 below gives the electron affinities of the main group of elements. If the ion that is formed by gaining an electron is stable, the electron affinity is expressed as a negative integer. The more ________________________ the ion, the higher is the negative integer for the electron affinity. Notice that fluorine has the highest electron affinity. This indicates that fluorine is very likely to be involved in chemical reactions. In fact, fluorine is the most ________________________of all elements. Metals have ___________ electron affinities. That is especially true for the Group 1 (1A) and 2 (2A) elements. Atoms of these elements form ________________________positive ions. A negative ion that is formed by elements of these groups is unstable. It breaks apart into a neutral atom and a free electron
Fig. 2.18
Text p. 57
Fig 2.19
p. 58 text
Periodic Trend for Electronegativity (p.p. 70 – 71) When two atoms form a bond, each atom attracts the other atom’s electrons in addition to its own. The electronegativity of an atom is a measure of the atom’s ability to ________________________ electrons in a chemical bond. _________ is used to symbolize electronegativity. There is a specific electronegativity associated with each element. While atomic radius, ionization energy, and electron affinity are all properties of single atoms, electronegativity is a property of atoms that are involved in chemical bonding.
Fig. 3.6
Text p. 71
The trend for electronegativity is the ________________________of the trend for atomic radius. In general, atomic size decreases from left to right across a period, electronegativity ________________________. This can be explained by looking at two factors. The number of ________________________ (which attract electrons) is increasing and at the same time, the number of energy levels (which shield the electrons from the attraction of the protons in the nucleus) remains the same. Thus electrons are pulled more tightly to the nucleus, resulting in a ________________________atomic size. In the second period, for example, lithium has the largest atomic radius and the lowest elecronegativity. As atomic size decreases across the second period, the electronegativity increases. Fluorine has the smallest atomic size in the third period, yet it has the largest electronegativity. Because noble gases do not usually participate in chemical bonding, their electronegativities are not given. Similarly, as atomic size increases down a group, electronegativity ________________________. As you move down a group, valence electrons are less strongly attracted to the nucleus because the number of energy levels between the nucleus and the valence electrons increases. In a compound, increasing energy levels between valence electrons and the nucleus mean that the nucleus attracts bonding pairs less strongly. For example, in Group 2 (IIA), berylliun has the smallest atomic radius and the largest electronegativity. As atomic size increases down a group, electronegativity ________________________. Section Review
1. How does your knowledge of electron arrangement and forces in atoms help you explain the following periodic trends ?
a. atomic radius b. ionization energy c. electron affinity
2. Using only their position on the periodic table, rank each of the following sets of elements in order of increasing atomic size. Explain your answer in each case.
a. Mg S Cl b. Al B In
c. Ne Ar Xe d. Rb Xe Te
e. P Na F f. O S N
3. Using only their position on the periodic table, rank each of the following sets of elements in order of decreasing ionization energy. Explain your answer in each case.
a. Cl Br I b. Ga Ge Se
c. K Ca Kr d. Na Li Cs
e. S Cl Br f. Cl Ar K
4. Which element in each of the following pairs will have the lower electron affinity ? Explain your answer.
a. K or Cl b. O or Li
c. S or Se d. Cs or F
Notes
Notes
Chapter 2 Review
p. 61
Notes
Chapter 2 Review
P. 62
Notes
Chapter 2 Review
P. 63
3.1
Classifying Chemical Compounds (p.p. 66 – 74) Most elements do not exist in nature in their pure form as ________________________. Gold, silver, and platinum are three metals that can be found in the earth’s crust as elements. They are called “precious metals” because their occurrence is so rare. Most other metals, and most other elements, are found in nature only as ________________________. There are only about 90 naturally occurring elements. In comparison, there are thousands upon thousands of different compounds in nature, and more are constantly being discovered. Because there are so many compounds, chemists have developed a classification system to organize them according to their properties, such as melting point, boiling point, hardness, conductivity and solubility.
Express Lab A Metal and a Compound Text p. 67
Complete the Express Lab along with Analysis Questions 1 – 4
Notes
Properties of Ionic and Covalent Compounds Based on their physical properties, compounds can be classified into two groups: ________________________compounds and ________________________compounds. Some of the properties of ionic and covalent compounds are summarized in the table below:
PROPERTY
IONIC COMPOUND
COVALENT COMPOOUNDS
State at room temperature
Melting Point
Electrical conductivity as a liquid
Yes
Solubility in water
Conducts electricity when dissolved in water
What is Bonding ? Chemical bonds are the forces that ________________________atoms to each other in compounds. Bonding involves the interaction between the ________________________ electrons of atoms. Usually the formation of a bond between two atoms creates a compound that is more ________________________than either of the two atoms on their own. The different properties of ionic and covalent compounds result from the manner in which chemical bonds form between atoms in these compounds. Atoms can either ________________________or ________________________electrons. When two atoms exchange electrons, one atom loses its valence electron(s) while the other atom gains the electron(s). This kind of bonding usually occurs between a metal and a non-metal. Recall, that metals have low ionization energies, while non-metals have high electron affinities. That is, metals tend to ________________________electrons, and non-metals tend to ________________________electrons. When atoms exchange electrons, they form an ________________________bond. Atoms can also share electrons. This kind of bond forms between two ________________________. It can also form between a metal and a non-metal when the metal has a high ionization energy. When atoms share electrons, they form a ________________________ bond. Each atom contributes one or more electrons to the covalent bond. The electrons are shared between the atoms.
Predicting Bond Type Using Electronegativity (p. 72 Text) You can use differences in electronegativities to decide whether the bond between atoms in ionic or covalent. When calculating electronegativity differences, the smaller value is always subtracted from the larger value so that the difference is always positive.
The Range of Electornegativity Differences When two atoms have electronegativity values that are identical, as in oxygen, they share the bonding their bonding pair of electrons equally between them in a pure covalent bond. When two atoms have electronegativity values that are different as in potassium fluoride, potassium gives up its valence electron to fluorine. Therefore, the bond is ________________________. It is not always clear whether atoms share electrons or transfer them. Atoms with different electronegativities can share electrons unequally without exchanging them. The figure below shows the range of electronegtivity differences. These values go from mostly covalent at 0.0 to mostly ionic at 3.3. Chemists consider bonds with an electornegativity difference that is greater than 1.7 to be ________________________and bonds with an electronegativity difference less than 1.7 to be ________________________.
Fig 3.8
p. 73 text
The table above shows you how to think of bonds having a percent ionic character, or a percent covalent character based on their electronegativity differences. When bonds have nearly 50 % ionic or covalent character, they have characteristics of both types of bonding. Covalent bonds between two atoms with different electronegativity values will have a certain degree of ionic character and are said to be ________________________covalent bonds.
Fig. 3.6 p. 71 text
Practice Problems
1. Determine the ΔEN for each bond shown. Indicate whether each bond I ionic or covalent
a. O - H
b. C - H
c. Mg - Cl
d. B - F
e. Cr - O
f. C - N
g. Na - I
h. Na - Br
Section Review
1. Name the typical properties of an ionic compound. Give two examples of ionic compounds.
2. Name the typical properties of a covalent compound. Give two examples of covalent compounds.
3. Describe and explain the periodic trend for electronegativity.
4. Based only on their position in the periodic table, arrange the elements in each set in increasing attraction for electrons in a bond.
5. Determine the ΔEN for each bond. Indicate whether the bond is ionic or covalent.
a. N - O b. Mn - O c. H - Cl d. Ca - Cl
Notes
Notes
3.2 IONIC AND COVALENT BONDINNG: THE OCTET RULE
The Octet Rule When atoms form bonds, they are often more stable. We know that noble gases are the most stable elements in the periodic table. They are extremely unreactive, and they tend not to form compounds. What the noble gases have in common is a _____________________ outer electron energy level. When an atom loses or gains electrons to achieve a filled outer electron energy level, the atom often becomes more stable. According to the octet rule, atoms bond in order to achieve an electron configuration that is the same as the electron configuration as a noble gas. When two atoms or ions have the same electron configuration, they are said to be isoelectronic. For example, Cl - is isoelctronic with Ar because both have 18 electrons and a filled outer energy level.
Ionic Bonding Sodium has a very low electornegativity, and chlorine has a high electronegativty. Therefore, when sodium and chlorine interact, sodium transfers its valence electron to chlorine. As shown in the figure below, sodium becomes Na + 1 and chlorine becomes Cl -.
Fig 3.9
p. 75 text
When neutral sodium loses its one valence electron to chlorine, the resulting Na + 1 cation has an en electron energy level that contains ___________ electrons. It is _____________________ with the noble gas neon. On the other hand, chlorine’s outer energy level has seven electrons. When chlorine gains sodium’s electron, it becomes an anion _____________________ with the noble gas argon.
The figure below shows how to represent the formation of an ionic bond using Lewis Structures. Thus, in an ionic bond, electrons are _____________________ from one atom to another so that they form oppositely charged ions. The strong force of attraction between oppositely charged ions is what holds them together.
Fig. 3.10
Text
Fig. 3.11 text Practice Problems
1. For each bond below, determine the ΔEN. Is the bond ionic or covalent ? a. Ca - O
b. K - Cl
c. K - F
d. Li - F
e. Li - Br
f. Ba - O
2. Draw Lewis structures to represent the formation of each bond in Question 1
Notes
COVALENT BONDING
We have just learned what happens when the electronegativity difference between two atoms is greater than 1.7. The atom with the lower electronegativity transfers its valence electron(s) to the atom with the higher electronegativity. The resulting ions have opposite charges. They are held together by a strong ionic bond. What happens if the electronegativity difference is zero. As an example, consider chlorine gas. Each chlorine atom has seven electrons in its outer energy level (valence shell). In order for chlorine to achieve the electron configuration of a noble gas according to the octet rule, it needs to gain one electron. When two chlorine atoms bond together, their electronegativity difference is zero and the electrons are equally attracted to each atom. Therefore, instead of transferring electrons, the two atoms each share one electron with each other. In other words, each atom contributes one electron to a covalent bond. A covalent bond consists of a pair of _____________________ electrons. Thus, each chlorine atom achieves a filled outer energy level, satisfying the _____________________ rule. The figure below shows how to represent a covalent bond with a Lewis structure.
Fig . 3.15 text
When two atoms of the same element form a bond, they share their electrons equally in a _____________________ covalent bond. Elements that bond to each other in this way are known as _____________________ molecules. When atoms such as carbon and hydrogen bond to each other, their electronegativities are so close that they share electrons almost equally. Carbon and hydrogen have an electronegativity difference of only 2.6 -= 2.2 = 0.4. In the figure below, you can see how one atom of hydrogen forms a covalent bond with four atoms of hydrogen. The compound methane, CH4 , is formed.
Fig 3.16 text
Each hydrogen shares one of its electrons with carbon. The carbon shares its four electrons with each hydrogen atom. Thus, each hydrogen achieves a filled outer energy level, as does carbon. (Recall that elements in the first period only need two electrons to fill their outer energy level). When analyzing Lewis structures that show covalent bonds, count the shared electrons as if they belong to each of the bonding atoms.
Practice Problems
1. Show the formation of a covalent bond between two atoms of each diatomic element. a. iodine
b. bromine
c. hydrogen
d. fluorine
2. Use Lewis structures to show the simplest way in which each pair of elements forms a covalent bond, according to the octet rule. a. hydrogen and oxygen
b. chlorine and oxygen
c. carbon and hydrogen
d. iodine and hydrogen
e. nitrogen and hydrogen
f. hydrogen and rubidium
Mutliple Covalent Bonds (p. 82 Text) In covalent bonding, atoms sometimes need to share two or three pairs of electrons, according to the octet rule. For example, consider the diatomic element oxygen. Each oxygen atom has six valence electrons in its outer shell, and therefore, each atom requires two additional electrons to achieve a stable octet. When two atoms form a bond by sharing two pairs of electrons, it results in a _____________________ covalent bond. When atoms form a bond by sharing three pairs of electrons, they form a _____________________ bond. An example of a triple bond is seen in diatomic nitrogen.
Figures 3.17, 3.18, 3.19 p. 82 Text
Note: the number of bonds each atom must form is equal to its ________________ Section Review
1. Use Lewis structures to show how each pair of elements forms an ionic bond.
a. magnesium and fluorine b. potassium and bromine
c. rubidium and chlorine d. calcium and oxygen
2. Use Lewis structures to show how each pair of elements forms covalent bonds.
a. one silicon atom and two oxygen atoms
b. one carbon atom, one hydrogen atom, and three chlorine atoms
c. two nitrogen atoms
d. two carbon atoms bonded together with three hydrogen atoms bonded to one carbon atom and one hydrogen atom and one oxygen atom bonded to the other carbon.
3. Use what you know about electronegativity differences to decide what kind of bond would form between each pair of elements.
4. In general, the farther away two elements are form each other on the periodic table, the more likely they are to participate in ionic bonding. Do you agree with this statement? Explain why or why not.
Notes
Polar Covalent Bonds: The “In-Between Bonds”
When two atoms have an elecronegativity difference hat is greater than 0.5 but less than 1.7, they are considered to be a polar covalent bond. In a polar covalent bond, the atoms have significantly different electronegativities. The electronegativity difference is not great enough, however, for the less electronegative atom to transfer its valence electrons to the other, more electronegative atom. The difference is great enough for the bonding electron pair to spend more time near the more electronegative atom. For example, the bond between oxygen and hydrogen in water has an electronegativity difference of 1.24. Because this value is less than 1.7 but greater than 0.5, the bond is a polar covalent bond. The oxygen attracts the electrons more strongly than the hydrogen. Therefore, oxygen has a slightly negative charge and the hydrogen has a slightly positive charge. These charges are not full charges but rather a partial charge symbolized by the symbols δ + and δ -.
Fig 3.24 & Fig 3.25
P. 86 Text
Practice Problems Predict whether each bond will be covalent, polar covalent, or ionic.
a. C – F
b. O – N
c. Cl – Cl
d. Cu – O
e. Si – H
f. Na – F
g. Fe – O
h. Mn – O
2. For each polar covalent bond in Question 1, indicate the location of the partial charges.
3. Arrange the bonds in each set in order of increasing polarity. (A completely polarized bond is an ionic bond)
a. H – Cl O – O Na – Cl
b. C – Cl Mg – Cl P – O N - N
Notes
Notes
Fig. 3.26
p. 87 text
Lab of some sort
Section Review
p. 94 text questions
Notes
Appendix 1 Periodicity and Bonding Questions
Review Perodicity & Bonding p. 1
Sch3a Worlbook
Appendix 1 Periodicity and Bonding Questions
Review Perodicity & Bonding p. 2
Sch3a Worlbook
3.4
WRITING CHEMICAL FORMULAS AND NAMING CHEMICAL COMPOUNDS
A chemical formula provides two important pieces of information:
• the elements that make up the compound
• the number of atoms of each element that are present in a compound
The order in which the elements are written also communicates important information. Convention, the less electronegative element is always listed first in the formula, the more electronegative element comes second. For example, the ionic compound that is formed from calcium and bromine is written CaBr 2. Calcium, a metal with a low electronegativity, is listed first. The subscript 2 after bromine indicates that there are 2 bromine atoms for every calcium atom.
Fig. 3.39 & 3.40
P. 95 Text
Then all of Unit 3 from Workbook, beginning with Nomenclature notes at the end of unit 3
CHAPTER 4: CLASSIFYING CHEMICAL REACTIONS: CHEMICALS IN BALANCE
4.1 Chemical Equations Any substance that undergoes a chemical reaction is called a _____________________. A substance that is formed is called a _____________________. Chemists use chemical equations to communicate what is happening in a chemical reaction. Chemical equations come in several forms. Word Equations
A word equation identifies the reactants and products by name. For example, sodium + chlorine → sodium chloride In this equation, “+” means “reacts with” and “→” means “to form”. Practice Problems
1. Describe each reaction using a word equation. Label the reactants and products.
a. calcium and fluorine react to form calcium fluoride
b. barium chloride and hydrogen sulphate react to form hydrogen chloride and barium sulphate
c. calcium carbonate, carbon dioxide and water react to form calcium hydrogen carbonate
d. hydrogen peroxide reacts to form water and oxygen
e. sulfur dioxide and oxygen react to form sulfur trioxide
NOTES
Word equations are useful because they identify the products and reactants in a chemical reaction. They do not, however, provide any chemical information about the compound or elements themselves. Another shortcoming of the word equation is that some names for chemicals are very long and awkward to work with. Skeletal Equations A skeletal equation lists the chemical formula of each reactant on the left, separated by a “+” sign if more than one reactant is involved, followed by an arrow “→”. The chemical formula for each product is listed on the right, again separated by a “+” sign if more than one product is formed. A skeletal equation shows the state of each reactant by using the appropriate subscript.
For example, the reaction of sodium metal with chlorine gas to form sodium chloride can be represented by the following skeletal equation:
Na (s) + Cl 2 (g) → NaCl (s)
A skeletal equation is more useful to a chemist than a word equation because it show the formulas of the compounds involved. It also shows the state of each substance.
SYMBOL
+
→
(s)
(l)
(g)
(aq)
MEANING
Practice Problems
20. Write a skeletal equation for each of the following:
a. solid zinc reacts with chlorine gas to form solid zinc chloride
b. solid calcium and liquid water react to form solid calcium hydroxide and hydrogen gas.
c. Solid barium reacts with solid sulfur to produce solid barium sulfide
d. Aqueous lead (II) nitrate and solid magnesium react to form aqueous magnesium nitrate and solid lead.
21. In each reaction below, a solid reacts with a gas to form a solid. Write a skeletal equation for each reaction.
a. carbon dioxide + oxygen → calcium carbonate
b. aluminum + oxygen → aluminum (III) oxide
c. magnesium + oxygen → magnesium oxide
NOTES
NOTES
Why Skeletal Equations Are Incomplete Although skeletal equations are useful, they do not fully describe chemical reactions. For example, according to the skeletal equation showing the formation of sodium chloride, one molecule of sodium reacts with one chlorine molecule containing two chlorine atoms. The product is one formula unit of sodium chloride containing one atom of sodium and one atom of chlorine. Where has the extra chlorine atom gone ?
The Law of Conservation of Mass All atoms must be accounted for according to the Law of Conservation of Mass. The Law of Conservation of Mass states that in any chemical reaction, the_____________________
______________________________________________________________________________. In other words, according to this law, matter cannot be _____________________ nor _____________________. Balanced Chemical Equations A balanced chemical equation reflects the Law of Conservation of Mass. This type of equation shows that there is the same number of each type of atom on both sides of the equation. Some skeletal equations are already balanced. For example, the skeletal equation for the reaction between carbon and oxygen to form carbon dioxide shows one carbon and two oxygen atoms on the reactant side of the equation, and one carbon and two oxygen atoms on the product side of the equation
Fig 4.1
p. 114 text
Most chemical equations, however, are not balanced, such as the one showing the formation of sodium chloride.
Fig. 4.2, 4.3, 4.4
p.p. 115 of text delete writing in between
You cannot balance an equation by changing any of the chemical formulas. The only way to balance an equation is to put the appropriate numerical coefficient in front of each compound or element in the equation.
Practice Problems
22. Copy the following skeleton equations and balance them.
a. S (s) + O 2 (g) → SO 2 (g)
b. P 4 (s) + O 2 (g) → P 4 O 10 (s)
c. H2 (g) + Cl 2 (g) → HCl (g)
d. SO 2 (g) + H2 O (l) → H2 SO4 (aq)
23. Indicate whether these equations are balanced. If they are not, balance them.
a. 4 Fe(s) + 3 O 2 (g) → 2Fe2O3(s)
b. HgO(s) → Hg (l) + O 2 (g)
c. H2 O2 → 2 H2 O (l) + O 2 (g)
d. 2 HCl (aq) + Na2SO3 (aq) → 2 NaCl (aq) + H2 O (l) + SO 2 (g)
NOTES
Steps for Balancing Equations
Here are some steps to follow when balancing equations that are more complex:
1. Balance the element, other than hydrogen and oxygen that has the greatest number of atoms in reactant or product.
2. Balance the other elements other than hydrogen and oxygen
3. Balance hydrogen and oxygen, whichever is present in the combined state. Leave until last, whichever one is in uncombined state
4. Check that each equation is balanced by counting the number of atoms of each element on each side of the equation
When the equation is balanced, the coefficients should be in their lowest terms.
Some balancing equations from SCH3a workbook 2000
p. 117 Text
p. 118 Text
4.2
SYNTHESIS AND DECOMPOSITION REACTIONS
(p.p. 119 – 125 Text) Just as there are different types of compounds, there are different types of chemical reactions. We will learn about ______ major classifications used for chemical reactions. Synthesis Reactions In a synthesis reaction, two or more elements or compounds combine to form a new substance. Synthesis reactions are also known as combination or formation reactions. A general equation for a synthesis reaction is In a simple synthesis reaction, one element reacts with one or more other elements to form a compound. There are several types of synthesis reactions and recognizing the patterns of the various types will help you predict whether substances will take part in a synthesis reaction. When a metal or a non-metal react with oxygen, the product is an _____________________. For example, the rusting of iron reacts according to the following equation:
A second type of synthesis reaction involves the reaction of a metal and a non-metal to form a binary compound. For example,
2 K (s) + Cl2 (g) → 2 KCl (s)
There are many types of synthesis reactions in which one or more compounds are the reactants. We will look at synthesis reactions of this type involving compounds oxides and water. When a non-metallic oxide reacts with water, the product is an _________________. The acids that form when a non-metallic oxide and water react are composed of hydrogen cations, and polyatomic anions containing oxygen and a non-metal. For example, one contributor to acid rain is sulfuric acid that is formed when sulfur trioxide reacts with water:
SO3 (g) + H2O(l) → H 2 SO4 (aq)
When a metallic oxide reacts with water, the product is a metal hydroxide, which is a ______________. For example, in the formation of calcium hydroxide from calcium oxide:
CaO (s) + 2 H2O (l) → Ca(OH) (2) (aq)
Practice Problems p. 122 Text
Notes
Notes
Decomposition Reactions
In a decomposition reaction, a compound breaks down into elements or other compounds. Therefore, a decomposition reaction is the ________________________of a synthesis reaction. The general formula for a decomposition reaction is:
_______________________________________
The substances produced by a decomposition reaction can be either elements or other compounds. For example:
2 H2O → 2 H2 + O2 or the explosive decomposition of ammonium nitrate:
NH4NO3 (s) → N2O(g) + 2 H2O (g)
Practice Problems p. 123
Notes
Notes
Combustion Reactions A complete combustion reaction is the reaction of a compound or element with _____ to form the most common oxides of the elements that make up the compound. Combustion reactions are usually accompanied by the production of light and heat. The complete combustion of any compound containing carbon, hydrogen, and oxygen (such as ethanol, C2H2OH) produces ________________________________and ________________________. Compounds that contain elements other than carbon also undergo complete combustion to form stable oxides. In the absence of sufficient oxygen, carbon containing compounds undergo incomplete combustion leading to the formation of carbon ________________________, ________, and water. Carbon monoxide is a deadly gas. You should always be sure that you have sufficient oxygen in your indoor environment for your gas furnace, gas stove, or fireplace.
Practice Problems p. 124
Notes
Notes
Section Review
p. 125 Text
Notes
Notes
4.3
SINGLE DISPLACEMENT AND DOUBLE DISPLACEMENT REACTIONS
(p.p. 126 – 141 Text) Single Displacement Reactioons In a single displacement reaction, one element in a compound is ________________________ (or replaced) by another element. Two general reactions represent two different types of single displacement reactions. One involves a metal replacing a metal cation in a compound; For example:
Or an example
The second type of single displacement reaction involves a non-metal, usually a halogen, replacing an anion in a compound as follows:
Or an example
Single Displacement Reactions and the Metal Activity Series
Single displacement reactions 1, ,2 3
p. 126 text When analyzing single displacement reactions, use the following guidelines:
• Treat hydrogen as a metal
• Treat acids such as HCl as ionic compounds of the form H + Cl or sulfuric acid
(H2SO4 as H + H + SO42- )
• Treat water as ionic with the formula H + (OH - )
Practice Problems p. 127
Notes
Through experiments, chemists have ranked the relative reactivity of metals, including hydrogen in acids and water), in an activity series. The reactive metals, such as potassium, are at the top of the activity series while the unreactive metals such as gold are at the bottom.
The Metal Activity Series The more reactive metals are at the ________________________of the activity series, the more stable ones are at the bottom. A reactive metal will displace any metal in a compound that is ________________________it in the activity series. Metals from lithium to sodium will replace hydrogen as a gas from water. Metals from magnesium to lead will displace hydrogen as a gas only form acids. Copper, mercury, silver, and gold will not displace hydrogen from acids. A single displacement reaction always favors the production of the ____________ reactive metal. In other words, the free metal that is formed from the compound must always be less reactive than the metal that displaced it. For example:
2 AgNO3 (aq) + Cu (s) → Cu(NO3)2 (aq) + 2 Ag (s)
Silver metal is more stable than copper metal. In other words, silver is below copper in the activity series.
Table 4.2
p. 130 text
Inv 4-A
p.p. 128
Inv 4-A
p.p. 129
You can use the activity series to help you predict the products of a reaction of a metal and a metal containing compound. For example,
Fe(s) + CuSO4 (aq) →
Iron is above copper in the activity series, meaning iron can replace copper and the reaction will proceed.
Fe(s) + CuSO4 (aq) → FeSO4 (aq) + Cu(s)
In the following reaction involving silver and calciu chloride,
Ag (s) + CaCl2 (aq) →
Silver is below calcium on the activity series, meaning that it is less reactive, therefore no reaction would take place.
Practice Problems
Top of p.131 text
NOTES
Single Displacement Reactions Involving Halogens Non-metals, typically halogens, can also take part in single displacement reactions. For example, molecular chlorine can replace bromine from KBr, an ionic compound, producing bromine and potassium chloride.
Cl2(g) + 2 KBr(aq) → 2 KCl(aq) + Br2(l)
The activity series for the halogens directly ________________________the position of the halogens in the periodic table. It can also be shown in the following way. ________________________ is the most reactive, and ________________________is the least reactive.
F > Cl > Br > I
The activity series for halogen can be used the same way the activity series for metals is used. If the element is above another in the activity series, it will replace it. For example,
F2 (g) + 2NaCl (aq) → 2NaF (aq) + Cl2 (aq)
If the element is below another in the activity series, it will not replace it and no reaction will occur. For example,
I2 + CaBr2 (aq) → no reaction
Practice problems
Bottom of
p. 131 text
NOTES
Notes
Double Displacement Reactions
A double displacement reaction involves the ________________________of cations between two ionic compounds, usually in aqueous (water) solution. A general equation for a double displacement reaction is:
____________________________________________
If you have 2 unlabelled beakers, one with distilled water, the other with salt water. You could determine the salt water solution by adding a few drops of silver nitrate solutions. The formation of a white precipitate (solid) indicates the presence of sodium chloride. The following double displacement reaction has occurred:
Fig . 4.15
p. 132 text
Since silver chloride is virtually insoluble in water, it forms a solid compound or precipitate. Double displacement reactions tend to occur in aqueous solutions. You can tell a double displacement reaction has taken place in the following cases:
• A solid (precipitate) forms
• A gas is produced
• Some double displacement reactions form a molecular compound like water. It is hard to tell when water has been formed, because often the reactions take place in water.
Double Displacement Reactions that Form a Precipitate A precipitate is a solid that separates from a solution as a result of a chemical reaction. Many double displacement reactions involve the formation of a precipitate. To predict whether a precipitate will form in a double displacement reaction, you must know the solubilities of different compounds. To determine the products, and their physical states in a double displacement reaction, you must first “deconstruct” the reactants, and then “reconstruct” the cations by switching the cations. You will be given the information on whether any of the products will form a precipitate.
Given the following reactants, how would you predict the products of the reaction and their state ? (Note, many hydroxide compounds are insoluble. Potassium cations form soluble substances with all anions)
MgCl2(aq) + KOH(aq) →
Fig. 4.16
p. 133 text If both products are soluble, and neither product precipitates out, both ionic compounds are dissolved in water and no reaction occurs.
Practice problems
p. 134 text
Notes
NOTES
Double Displacement Reactions that Produce a Gas In certain cases, you know that a double displacement reaction has occurred because a gas is produced. The gas is formed when one of the products of the double displacement reaction decomposes to give water and a gas. For example, the reaction between sodium carbonate and hydrochloric acid. If you carry out this reaction, you immediately see the formation of carbon dioxide gas. The first reaction that takes place is a double displacement reaction. Determine the products in the following way.
• Separate the reaction into ions, and switch the anions. Write the chemical formulas for the products and balance the equation.
Na2CO3 (aq) + 2HCl (aq) → 2NaCl (aq) + H2CO3 (aq)
The carbonic acid is unstable and decomposes to carbon dioxide and water.
H2CO3 (aq) → H2O (l) + CO2 (g
Practice Problems
p. 135 text top of page
Notes
The Formation of Water In Neutralization Neutralization reactions are a special type of double displacement reaction that produces water. Neutralization involves the reaction of an acid with a base to form water and an ionic compound. For example, the neutralization of nitric acid with sodium hydroxide (a base) is a double displacement reaction.
HNO3 (aq) + NaOH (aq) → NaNO3 (aq) + H20 (l)
Often, neutralization reactions produce no precipitate or gas.
practice problems bottom p. 135 text
Notes
Investigation 4-B p. 146 text
Inv 4-B p. 137 text
Inv. 4-c p. 138 Tetx
Inv 4-v p. 139 text
Concept Organizer
p. 140 text Section Review Questions
p. 140 text
Section Review Questions
p. 141 text
Notes
Notes
Chapter 4 Review p. 149 text
Ch 4 Review
p. 150
NOTES
NOTES
Unit 1 review
p. 154 text
Unit 1 review
p. 155 text
Unit 1 review
p. 156 text
Unit 1 review
p. 157 text
Notes
Notes
Notes
Notes
Date: Name: Class:
Problem Solving STSE Connections in the Great Lakes
Chapter 1
BLM 1-1
Goal
Outline the connections among science, technology, society, and the environment in relation to the Great Lakes.
Procedure
1. Choose an Area of Concern or a more general issue from those listed below. Use print and electronic resources to investigate the chemicals involved, and the connections between science, technology, society, and the environment.
Great Lakes Areas of Concern
• Hazardous waste sites along the St. Lawrence river in Cornwall, Ontario, and in the Port Hope harbour, remain from past industrial waste handling practices.
• Sewage treatment plants and sewage overflow contributes to the pollution in the Bay of Quinte.
• Copper mining, milling, and smelting operations have contaminated the water in Torch Lake, connected to Lake Superior.
• Industrial chemical discharges pollute the Niagara River connecting Lake Erie and Lake Ontario.
• Urban run-off from Metro Toronto contaminates the water in Lake Ontario.
General Pollution Issues involving the Great Lakes
• The atmospheric/airborne deposition of toxic chemicals may have a large role in the pollution of the Great Lakes.
• Biomagnification and bioaccumulation of toxic chemicals through the food chain puts predators such as lake trout, salmon, herring gulls, and fish-eating humans at risk. (Hint: Check out fish consumption advisories for the Great Lakes.)
• Nuclear power plants discharge heated water into the Great Lakes. How does this affect the environment? Where is this happening?
2. Organize and present your findings using one or more media such as an essay, a web site, a PowerpointTM or CorelTM presentation, or a brochure.
Analysis
1. What chemicals are involved in the issue you investigated?
2. For what purpose were the chemicals originally produced?
3. How did the chemicals benefit society?
4. How did the chemicals affect the environment?
5. What is being done to solve the problem?
Conclusions
6. Use a concept web to illustrate the STSE connections present in this issue.
UNIT 2
CHEMICAL
REACTIONS
UNIT 1
MATTER AND CHEMICAL BONDING
UNIT 1
MATTER AND CHEMICAL BONDING
1.1 The Study of Chemistry
What is chemistry ?
1.2
Define matter:
What are properties?
What is a physical property?
A physical property is s property that you can observe without changing one kind of matter into something new.
What is a chemical property?
Using Measurement to Describe Matter
What is the name of the system of measurement that scientists rely on to communicate effectively? What does it allow scientists to do?
Name five of the base S.I. units?
Complete the table below:
Table 1: Important SI Quantities and Their Units Quantity Definition S.I. Units or their derived equivalents Equipment used to measure the quantity
mass
length
temperature
volume
Mole
Density
Energy
Measurement and Uncertainty
What are exact numbers?
What two factors affect your ability to communicate measurements and calculations?
Significant Digits, Certainty, and Measurements
Calculated answers have the same number of significant digits as the least-certain measurement involved in the calculation; digits in excess of the number of least-certain digits must be rounded off when expressing the final answer.
The digits that you record when you measure something are called significant digits. Significant digits include all the digits that you are certain about and a final, uncertain digit that you estimate.
How Can You Tell which Digits are Significant ?
Complete the following table of rules for determining significant digits:
Rules Examples
Explaining Three Significant Digits in the Example Above
Explaining Five Significant Digits in the Example Above
Complete Practice Problems #1, 2 (text p. 18)
What is Precision?
Complete Investigation 1-A Observing Aluminum Foil (text p. 13)
Calculating With Significant Digits (text p. 20 – 22)
In this course, you will often take measurements and use them to calculate other quantities. You must be careful to keep track of which digits are significant.
Why is it important to keep track of significant digits?
Complete the rules for reporting significant digits in calculations in the table below:
Table 1 Rules for Reporting Significant Digits in Calculations
Rule 1 Multiplying and Dividing
Rule 2 Adding and Subtracting
Rule 3 Rounding
Sample Problem Reporting Volume Using Significant Digits
A student measured a regularly shaped sample o f iron and found it to be 6.78 cm long, 3.906 cm wide and 11 cm tall. Determine its volume to the correct number of significant digits.
Length = 6.78 cm (3 sig digits)
Width = 3.906 cm (4 sig digits)
Height = 11 cm (2 sig digits)
Solution
Sample Problem Reporting Mass Using Significant Digits
Suppose that you measure the masses of four objects as 12.5 g, 145.67 g, 79.0 g, and 38.438 g.
What is the total mass of the four objects ?
Strategy:
Complete Practice Problems # 3 a – g, p. 22 text) Reinforcement Reporting Significant Digits in Calculations
Chapter 1
BLM 1-2
Goal
Reinforce your understanding of significant digits.
Procedure
1. State the number of significant figures in each of the following:
(a) 3570
(b) 17.505
(c) 41.400
(d) 0.51
(e) 0.000 572
(f) 0.009 00
(g) 41.50 × 10-4
(h) 0.007 160 × 105
(i) 1.234 00 × 108
(j) 0.000 410 0 × 107
2. Perform the following operations and give the answer to the correct number of significant digits.
(a) 15.1 + 75.32
(b) 178.904 56 – 125.805
(c) 4.55 × 10-5 – 3.1 x 10-5
(d) 0.000 159 + 4.0074
(e) 1.805 × 104 + 5.89 × 104
(f) 0.000 817 - 0.000 048 1
(g) 8.166 × 105 – 7.819 × 105
(h) 45.128 + 8.501 87 – 42.18
(i) 5.677 × 10-6 + 7.785 × 10-6
(j) 8.75 × 10-9 + 6.1157 × 10-9
(k) 1.99 ÷ 3.1
(l) 1200.0 ÷ 3.0
(m) 5.32 x 10-4 ÷ 4.218 × 10-8
(n) 45.32 x 2.3
(o) 0.024 00 ÷ 6.000
(p) 12.4 x 0.30
(q) (5.50 x 108) × (4 x 105)
(r) 7.4 ÷ 3
(s) 4.75 × 5
(t) 2.5 × 6.700 ÷ 0.891
Date: Name: Class:
Assessment
Chapter 1
BLM 1-4
Goal
Demonstrate and assess your understanding of the concepts presented in Chapter 1
Part 1:
Multiple
Choice
Circle the letter corresponding to the best answer.
1. The metric base unit to measure linear dimensions is the:
(a) kelvin
(b) metre
(c) candela
(d) kilogram
(e) second
2. Which of the following is not an SI base unit?
(a) litre
(b) second
(c) ampere
(d) kilogram
(e) mol
3. Which digit or digits are uncertain in the result of a scientific measurement?
(a) the last
(b) the last two
(c) all of them
(d) none of them
(e) the first
4. How many significant digits are contained in the number 0.000540?
(a) two
(b) three
(c) five
(d) six
(e) seven
5. Which of the following is characteristic of pure substances?
(a) They are solutions.
(b) They are mixtures.
(c) They are composed of one element.
(d) They have a fixed chemical composition.
(e) They have variable composition.
6. All mixtures:
(a) are easy to separate
(b) have a constant boiling point
(c) are composed of parts with different identities
(d) have a chemical formula
(e) are the result of chemical change
7. An example of a pure substance is:
(a) salt water
(b) copper(II) chloride
(c) air
(d) tap water
(e) soda water
Part 2: Short Answer & Calculations 8. Carry out the following metric conversions.
(a) 45.6 g to mg
(b) 0.00867 kL to mL
9. Perform the following calculations and express the answer to the correct number of significant digits.
(a) 0.000354 × 2.8
(b) 7.90 + 3.222
(c) 8.432 ÷ 2.5
(d) 0.17 – 0.154
10. Identify the following as being either a physical or a chemical change.
(a) a chicken cooking in the oven
(b) ice cubes subliming
(c) iron rusting in air
(d) water evaporating form salt water in a beaker when heated
(e) toasting a slice of bread
(f) burning a log
11. Round off the following to the number of significant digits indicated.
(a) 0.0002554, to 2 significant digits
(b) 35.8348, to 4 significant digits
(c) 11.4555, to 3 significant digits
12. How long is the object shown below? Record your answer using the correct number of significant digits.
1.3 Classifying Matter and Its Changes Matter changes in response to changes in energy.
What is energy?
What is a physical change?
What is a chemical change
Classifying Matter
What is a mixture ?
What is a heterogeneous mixture?
What is a homogeneous mixture? What’s another name given?
What is a pure substance?
How are pure substance further classified?
What is an element? Give some examples.
What is a compound? Give some examples.
Chapter 1 Review (text p.p. 29 – 31) #2, 5, 6, 7, 8, 9, 10, 11, 12,
Chapter 2 Elements of the Periodic Table (text p. 33 – 64)
2.1 ATOMS AND THEIR COMPOSITION (text p.p. 34 – 39)
THE MODERN VIEW OF THE ATOM
What is the smallest unit of an element that still retains the properties of that element?
Atoms themselves are made of even smaller particles. Name the three subatomic particles and give their location in the atom.
THE ATOMIC THEORY OF MATTER
Who was John Dalton and what is he famous for? What was his model called?
Describe the four main points of Dalton’s atomic theory.
Dalton’s concept of the atom was one of several ideas in his atomic theory of matter. Dalton’s atomic theory of matter (1809) has four main points:
•
• •
•
• Law of Conservation of Mass:
• Law of Definite Proportions:
Complete the table below:
Table 2.1 Properties of Protons, Neutrons, and Electrons
Subatomic Particle
Charge
Symbol
Mass (g)
Radius (m)
Electron
Proton
Neutron
EXPRESSING THE MASS OF SUBATOMIC PARTICLES • What is an Atomic Mass Unit
The Nucleus of an Atom
• What does Atomic Number represents? ( Z )
• What does the Mass Number represents? ( A )
Information about an element’s protons and neutrons is often summarized using the chemical notation shown in Figure 2.3.
Mass # A
X
Atomic # Z
In the figure above, which letter represents the atomic symbol of the element ?
In the following example, identify the mass number and the atomic number
Atomic number
Mass number
19 F 9
For the isotope of fluorine above, what do the numbers represent?
How can you determine the number of neutrons in the nucleus of an atom?
How many neutrons would this isotope of fluorine have? PRACTICE PROBLEMS
1. Complete the missing values or chemical notations about the atoms from the table on p. 37 of your textbook below.
a. b. c.
d. e f.
g h i
j k l
m n o
p. q. r.
s. t. u.
v. w. x.
When chemists refer symbolically to an element, they will often leave out the atomic number. For example, for oxygen, they would write 16 O or O – 16.
USING THE ATOMIC NUMBER TO INFER THE NUMBER OF ELECTRONS
How can you infer the number of electrons from the atomic number? Why can you do this?
ISOTOPES AND ATOMIC MASS
What is am isotope?
Compare the chemical properties of isotopes of the same element. Explain why it is like this.
What is radioactivity”?
Atomic Weights
What information does the average atomic weight of an element give? What other name can it have?
Write the formula used to calculate the average atomic weight.
Example 1 Sulfur has three isotopes which have the following relative abundances: S – 32 = 95% S – 33 = 0.8 % S – 34 = 4.2 % Calculate the atomic weight of sulfur.
ELECTRONS IN ATOMS
What subatomic particle is responsible for the chemical properties of an element?
REVISITNG THE ATOMIC THEORY
John Dalton did not even know about subatomic particles when he developed his atomic theory. Even so, the modern atomic theory, shown below, retains many of Dalton’s ideas, with only a few modifications.
The Modern Atomic Theory
SECTION REVIEW
1. For the table below, fill in the missing information in the appropriate space provided.
ELEMENT ATOMIC NUNMBER
MASS NUMBER
NUMBER OF PROTONS
NUMBER
OF ELECTRONS
NUMBER OF NEUTRONS
(a)
(b)
108
(c)
47
(d)
(e)
(f)
(g)
33
(h)
42
(i)
35
(j)
(k)
(l)
45
(m)
79
179
(n)
(o)
(p)
(q)
(r)
(s)
(t)
50
69
1. a. b. c. d.
e f. g. h.
i. j. k. l.
m. n. o. p.
q. r. s. t.
2. Explain the difference between an isotope and a radioisotope. Provide an example other than oxygen to support your answer.
3. Examine the information presented in the following pairs
3 3 14 16 19 18 H and He C and N F and F 1 2 6 7 9 9
(a) For each pair, do both members have the same number of protons ? electrons ? neutrons?
(b) Which pair or pairs consist of atoms that have the same value for Z? Which consists of atoms that have the same value for A ?
4. Compare Dalton’s atomic theory with the modern atomic theory. Explain why scientists modified Dalton’s theory.
5. Fill in the following table:
Protons
Electrons
Neutrons
P – 31
Al – 27
H-1 (protium)
H – 2 (deuterium)
EXPANDING THE MODEL OF THE ATOM (p.p. 656 – 659)
Emission Spectra
What can each wavelength of visible light be re[resented by?
What is a bright-line spectrum ?
Bohr’s Model of the Atom In 1913, Danish physicist, Niels Bohr developed a model of the atom that explained hydrogen emission spectrum.
Describe the location and motion of electros in Bohr’s a toms
The following three points of Bohr’s theory help explain hydrogen’s emission spectrum:
• In Bohr’s model, atoms have specific allowable ______levels. He called these energy levels stationary states. Each of these levels corresponds to a fixed, circular orbit around the nucleus.
• An atom does not give off energy while its electrons are in a stationary state
• An atom changes states by giving off, or absorbing a quantity of light energy exactly equal to the ______in energy between the two stationary states.
Bohr’s model was revolutionary because he proposed that the energy absorbed or emitted by an atom needed to have ______ ______. The energy change was ______, rather than continuous. When something is quantized, it is limited to discrete amounts or multiples of discrete amounts. Two great scientists paved the way for this surprising idea. German physicist Max Plank had already proposed that light was quantized, meaning that it exists in packets. Building on this idea, Einstein proposed that light could behave as particles, which he called photons. The energy associated with the light in a line spectrum corresponds with the change in energy of an electron as it moves _______or _______________ an energy level. For example, when electrons in hydrogen atoms that have been excited to the third energy level and subsequently drop to the second energy level, they emit light that has a specific wavelength. They emit photons of _____________ light that have a wavelength of 656.5 nm. These photons cause the red line spectrum for hydrogen seen in Figure D1 above.
Why was a new model required?
Sublevels The spectra of many electron atoms suggested that a more complex structure was needed. Notice that in Figure D1, the spectra for these more complex atoms have groups of lines close together. The groups of lines are separated by spaces. The large spaces represent energy differences between ______ ______, while the smaller spaces represent energy ______within levels. If the electrons are changing energy within the levels, this suggests that there are sublevels within each level., each with its own slightly different energy. The idea of different energy levels 1, 2, 3, 4, . . . remains, but each energy level is split up into sublevels called _____, _____, _____, _____. Examine the table below to see how electrons are arranged in sublevels for energy levels 1 to 4.
Each sublevel has its own energy. Examine the figure on the next page to see the relative energies of the different levels and sublevels. But where would you find electrons in this new model? How are the sublevels oriented in space?
Particles with Wave Like Properties
In 1924, Louis de Broglie suggested that all matter had -like properties. De Broglie developed an equation that allowed him to calculate the wavelength associated with any object. Objects we can see and interact with have calculated wavelengths that are smaller than electrons. Their wavelengths are so tiny compared to their size, that they do not have any measurable effect on the motion of the objects
For very tiny moving particles such as electrons, however, the wavelength becomes very significant. De Broglie’s theory was proven by experiment when streams of electrons produced diffraction patterns similar to electromagnetic radiation that was already known to travel in waves.
Orbitals
In 1920, Erwin Schrodinger used de Broglie’s idea that matter has wavelike properties. Schrodinger proposed the quantum mechanical model of the atom. In this new model, he abandoned the notion of the electron as a small particle orbiting the nucleus. Instead, he described the behaviour of electrons in terms of wave functions.
This imprecise nature of Schrodinger’s model was supported shortly afterwards by principle proposed by Heisenburg. Heisneburg demonstrated that it is impossible to know both an electron’s path and its exact location. Therefore it is impossible to know the exact location of an electron at any time, without some degree of uncertainty. Heisenburg’s uncertainty principle shows that you can never know both the ______and the ______ of an electron beyond a certain measure of precision. Heisneburg also showed that if you could know either the velocity or position precisely, then the other property would be uncertain. Therefore, when talking about where electrons are found, you can only talk in terms of
______ ______.
Schrodinger used a mathematical wave equation to define the probability of finding an electron within an atom. There are multiple solutions to this wave equation, and Schrodinger called these solutions ______. Each solution provides information about the energy and location of an electron within an atom. Each orbital has a specific energy associated with it, and each contains information about where inside the atom, the electrons would spend most of their time.
These regions of space where an electron is most likely to be found are called ______. In solving the wave equation, three unknowns were encountered (a fourth was added later). These unknowns are called the ______ ______and describe different properties of the electron.
The Four Quantum Numbers
1 . The Principle Quantum Number (_____)
• Indicates the ______an electron has
• Tells us which energy ______the electron is in
• Can have values of n = ______
• The higher the ‘n’ value, the ______the energy
• Correspond to Bohr’s orbits
2. The ______Quantum Number (_____)
• Indicates the ______of the orbital
• It can have values of L = ______ ______
• n- value possible ‘L’ values
1
2
3
4.
Note: the total number of possible ‘L’ values is the same as the ‘n’ value L- VALUE
0
1
2
3
SHAPE
SPEHRICAL
DUMBBELL
DOUBLE DUMBBELL
??
??
NAME
____– orbital
____– orbital
____– orbital
____– orbital
____– orbital
3. The Magnetic Quantum Number (M L)
• indicates the orientation of the orbitals about the nucleus.
• The number of values indicates the number of orbitals that exist with a given
• M L can have the values of - L . . . 0 . . . + L
L
0
1
2
3
4
5
NAME
M L
NUMBER OF POSSIBLE ORBITALS
4. _________________Quantum Number (_____)
• Related to the _________________ of an electron
• All electrons spin on their axis. The two different spins are assigned the value of
________________________
• indicates that each orbital can hold a maximum of ________ electrons. If there are two electrons in an orbital, they must have opposite spins.
Complete the following summary table on the Quantum Theory below:
n
L
NAME
M L # of
Orbitals
M S # of
Electrons Total # of
Electrons
1
2
2
3
3
3
4
4
4
4
The ‘n’ Rules
• The nimber of different shapes (or ) of orbitals in a given energy level =
• The total number of orbitals in a given energy level =
• The total number of electrons in a given energy level =
Shapes of Orbitals The shapes of the probability graphs of Schrodinger’s wave functions are the shapes of the orbitals in which electrons reside in atoms. You can visualize orbitals as electron clouds. The shape of each cloud is based on probability – it tells where the electron spends most of its time. Examine the figure below. The ‘s’ orbitals are spherical in shape. Each of these spherical shells contains two electrons. There are three ‘p’ orbitals for each sublevel, each with a capacity for two electrons. Each ‘p’ orbital is shaped something like a dumbbell. For each subleve, the ‘p’ orbitals are oriented along the x, y, and z axes. The d’ and ‘f’ orbitals are quite complex in shape. Each ‘d’ sublevel contains 5 ‘d’ orbitals and each ‘f’ sublevel contains 7 ‘f’ orbitals.
Filling Orbitals
What is the maximum number of electrons that can occupy any one orbital?
State the Pauli Exclusion Formula
How do electrons fill orbitals within atoms? They do so to minimize the potential energy of the atom.
1. The maximum number of electrons that any orbital can hold is _____________
2. They will fill the ________________________energy level first. The 1s orbital will fill before the 2s orbital, which fills before the 2 p orbitals and so on.
1. When occupying two or more orbitals with the same energy (for example any of the 3 ‘p’ orbitals), electrons will half fill each orbital until all are half filled before adding a second electron to each one. This is called ________________________Rule. You can show how electrons fill orbitals using superscripts in a notation called the electron configuration of the atom. For example, a boron atom contains 5 electrons. The electron configuration for Boron would be 1s 2 2s 2 2p 1. The electron configuration for Nitrogen is 1s 2 2s 2 2p 3.
Electron Configurations
Write the electron configuration for the following elements:
1. 21 Sc
2. 23 V
3. 30 Zn
4. 35 Br
5. 50 Sn
6. 57 La
7. 56 Ce
8. 64 Gd
9. 50 Hg
10. 84 Po
Quantum Review 1
ATOMS, ELEMENTS, AND THE PERIODIC TABLE (Text p.p. 40 – 47) By the mid 1800’s, there were 65 known elements. Chemists studied these elements intensively and recorded detailed information about their reactivity and the masses of their atoms. Some chemists began to recognize patterns and behaviour of many of these elements.
What was the contribution of Dmitri Mendeleev to the science of chemistry?
What was one of the major differences between Mendeleev’d original periodic table ad the new one?
State the Periodic Law
PRACTICE PROBLEMS
2. Identify the name and symbol of the elements in the following location of the periodic table:
a. Group 14 (IVA), Period 2
b. Group 11 (IB), Period 4
c. Group 18 (VIIIA), Period 6
d. Group 1 (IA), Period 3
e. Group 12 (IIB), Period 5
f. Group 2 (IIA), Period 4
g. Group 17, (VIIA), Period 5
h. Group 13 (IIA), Period 3
ELECTRONS AND THE PERIODIC TABLE
We have seen how the periodic table organizes elements so that those with similar properties are in the same group. You have also seen how the periodic table shows a clear distinction among metals, non-metals, and metalloids. Other details of the organization of the periodic table may seem strange. Why, for example, are there different numbers of elements in the periods.
The reason for this, and other details of the periodic table’s organization, involves the number and arrangement of ________________________in atoms of each element.
ELECTRONS AND ENERGY LEVELS Electrons cannot move haphazardly. Their movement around an atomic nucleus is restricted to fixed regions of space. These regions of space are three-dimensional. The following figure shows a representation of these regions.
These regions are not solid, but rather are volumes of space in which electrons may be found. They are referred to as energy shells or energy levels. An electron that is moving in a lower energy level is found ________________________to the nucleus. It would have less energy than it would if it were moving in a higher energy level. There is a limit to the number of electrons that can occupy each energy level. For example, a maximum of _______________ electrons can occupy the first energy level. A maximum of ________________________electrons can occupy the second energy level. The periodic trends that result from organizing the elements by their atomic number are linked to the way in which electrons occupy and fill energy levels.
As shown in the figure above, a common way to show the arrangement of electron in atoms is to draw circles around the atomic symbol. Each circle represents an energy level. Dots represent electrons that occupy each energy level. This kind of diagram is called a Bohr –Rutherford diagram. Figure 2.9 B shows that the first energy level is full when _____________ electrons occupy it. Only two elements have two or fewer electrons: hydrogen and helium. Hydrogen has one electron, and helium has two. These elements, with their electrons in the first energy level, make up Period _______ of the periodic table. As you can see in Fig. 2.9 C, Period 2 elements have two energy levels. The second energy level is full when ________electrons occupy it. Neon, with a total of ten electrons, has its first and second energy levels filled. Notice how the second energy level fills with electrons as you move across the period from lithium to fluorine. Figure 2.9 C
PATTERNS BASED ON ENERGY LEVELS AND ELECTRON ARRANGEMENTS The structure of the periodic table is closely related to ________________________ levels and the arrangement of electrons. Two important patterns result from this relationship. One involves periods, and the other involves groups. The Period Related Pattern
Elements in Period 1 have electrons in how many energy levels?
Elements in Period 2 have electrons in how many energy levels?
How many energy levels could an element in Period 5 occupy?\
The Group Related Pattern The second pattern emerges when you consider the electron arrangements in the main- group elements: the elements in Group 1 (1A), 2 (2A) and 13 (3A) to 18 (8A). All the elements in each main group have the same number of electrons in their highest (outer) energy level. What is the name given to electrons that occupy the outermost energy levels of an atom?
What in the atom is responsible for the chemical behaviour of an element?
USING LEWIS STRUCTURES TO REPRESENT VALENCE ELECTRONS It is time consuming to draw electron arrangements using Bohr-Rughterford diagrams. It is much simpler to draw Lewis structures to represent elements and their valence electrons. To draw a Lewis structure, you replace the nucleus and the inner energy levels of an atom with its atomic symbol. Then you place dots around the atomic symbol to represent the valence electrons. The order you place the first four dots is up to you, however, place a dot on each side of the symbol before doubling up on any dot.
Practice Problems
2. On the blank periodic table, sketch the electron arrangements for the first 20 elements using Lewis structures.
3. Use the periodic table to draw Lewis structures for the following elements: barium (Ba), gallium (Ga), tin (Sn), bismuth (B), iodine (I), cesium (Cs), krypton (Kr) and xenon (Xe)
4. Identify the number of valence electrons in the outer energy levels of the following elements:
a. chlorine f. lead
b. helium g. antimony
c. indium h. selenium
d. strontium i. arsenic
e. rubidium j. xenon
THE SIGNIFICANCE OF A FULL OUTER ENERGY LEVEL Describe the electron arrangement of the noble gases. What property does this give to the nobe gases?
How many valence electrons to the noble gases have?
What is a stable octet?
Summary
You have seen that the structure of the periodic table is directly related to energy levels and arrangement of electrons. The patterns that emerge from this relationship enable you to predict the number of valence electrons for any main group element. They also enable you to predict the number of energy levels that an element’s electrons occupy.
Orbital Box Notation Using this method, each orbital is shown as a box with ________________________ representing the electrons. If there are two electrons in the same orbital box, then they must have opposite spins. This is shown by having the arrows facing in different _____________________.
Element Electron Configuration Orbital Box
4 Be
6 C
7 N
Hund’s Rule _______________________________________________________________________________________________________________________
8 O
9 F
21 Sc
50 Sn
What is the ground state electron configuration of an atom?
Excited State Configurations In order to obtain these configurations, energy must be added to the atom. When energy is added, one or more electrons is moved to a higher energy orbital. To which orbital the electron gets moved depends on the amount of energy that is added.
Write the ground state and a possible excited state electron configuration for the following isotope of magnesium.
12 Mg Ground State Electron Configuration Excited State (one of many)
Electron Configurations and Valences
Why do atoms react with other atoms?
What is it in the atom that determines its chemical properties?
How can an atom become more stable?
Describe several stable arrangements of electrons in atoms.
Stable Arrangements atoms include:
1.
2.
3.
4.
Valences Predict the most likely valences for the following:
1. 7 N
2. 16 S
3. 30 Zn
4. 50 Sn
Explain why:
i. Fe (# 26) has a (+2) and a (+3) valence. Which is more stable ?
ii. Rh (# 45) has a (+2) and a (+4) valence.
iii. Cerium (# 58) has a (+3 and a (+4) valence
Exceptions to the Predicted Electronic Configurations There are two families whose electron configurations are different than what would be predicted. They are the families VIB and IB. The elements in these families spontaneously rearrange their electrons to make themselves more stable. To do this, they promote an electron from an ‘s’ orbital and excite it to a ‘d’ orbital.
Examples
24 Cr
42 Mo29
29Cu
47 Ag
Section Review
1. State the Periodic Law and provide at least two examples to illustrate its meaning.
2. Identify the Group number for each of these sets of elements. Then write the symbols for the elements within it:
• alkali metals
• noble gases
• halogens
• alkaline earth metals
3. Identify the elements that is described by the following information. Refer to a periodic table as necessary.
• It is a Group 14 (IIIA) metalloid in the third period
• It is a Group 15 (VA) metalloid in the fifth period
• It is the other metalloid in Group 15
• It is a halogen that exists in the liquid state at room temperature
4. What is the relationship between electron arrangement and the organization of the elements of the periodic table ?
5 How many valence electrons are there in an atom of each of these elements ?
(a) neon sodium magnesium bromine chlorine silicon sulfur helium strontium tin
(b) Present your answers from Part A in the form of Lewis structures.
(c) Classify each element from Part A as a metal, non-metal, or a metalloid
6. Compare and contrast noble gases with other elements.
7.
(a) Draw Lewis Structures for each of the following elements: lithium potassium magnesium aluminum carbon
(b) Which of these elements have the same number of energy levels ?
(c) Which have the same number of valence electrons
8. Identify the elements with the following electron distributions by orbits:
a. 2, 8, 8 b. 2, 8, 6 c. 2, 4 d. 2, 8, 8, 2
5. What is meant by a stable octet ?
6. How many atoms must each of the following elements gain to achieve a stable octet ?
a. fluorine b. oxygen c. nitrogen d. carbon
7. How many electrons must the following atoms lose in order to achieve a stable octet
a. silicon b. phosphorus c. sodium d. argon
8. How does the sign of the valence relate to the gain or loss of electrons ?
9. Explain why atoms form ions.
2.3
PERIODIC TRENDS INVOLVING THE SIZES AND ENERGY LEVELS OF ATOMS In Section 2.1, we learned that the size of a typical atom is about 10 – 10 m, however, we also know that the atoms of each element are distinctly different. For example, the atoms of different elements have different numbers of protons. This means of course, that they also have different numbers of electrons. You may predict that the size of an atom is related to the number of protons and electrons it has. Is there any evidence to support this prediction, If so, is there a pattern that could help you to predict the relative size of an atom for any element on the periodic table ? Chemists define, and measure, an atom’s size in terms of its radius. The radius of an atom is the distance from its nucleus to the approximate outer boundary of the cloud-like region of its electrons. This boundary is approximate because atoms are not solid spheres. They do not have an outer boundary.
Trends for Atomic Size (Radius) There are two general trends for atomic size, describe them.
•
•
Practice Problems
7. Using only their location in the periodic table, rank the atoms in each set by decreasing atomic size. Explain your rankings. (a) Mg Be Ba
(b) Ca Se Ga
(c) Br Rb Kr
(d) Se Br Ca
(e) Ba Sr Cs
(f) Se Br Cl
(g) Mg Ca Li
(h) Sr Te Se
(i) In Br I
(j) S Se O
EXPERIMENT THE REACTIVITY OF ELEMENTS
Trends for Ionization Energy
What is ionization energy?
What is an ion?
Name two different types of ions formed by the atom. Explain how the two different types of ions are formed?
The figure below shows the formation of ions for several elements. As you examine the diagrams, pay close attention to
• the energy level from which the electrons are gained or given up
• the charge of the ion that is formed when an atom gains or gives up electrons
• the arrangement of the electrons that remain after the electrons are gained or given up. If you try to visualize the periodic table as a cylinder, (see Fig. 2.15 below) rather than a flat plane, you can begin to see the relationship between ion formation and the electron arrangement of the noble gases.
Describe how metals and non-metals of the main group tend to react. What are they trying to become like?
Give an example.
How do elements in Group 1A react?
How do elements in Group 1A react?
How is the charge on a cation related to its group number?
Describe a genera rule for how non-metals react.
The figure below shows the ionization energy that is required to remove one electron from the outer energy level of atoms of the main group elements. This energy is called the ____________-ionization energy. It is measured in units of kJ / mol. As you can see, atoms that give up electrons easily have _______________ ionization energies. You would probably predict that the alkali metals of Group 1 (1A) would have low ionization energies. These elements are, in fact, extremely reactive because it takes so little energy to remove their single valence electron.
All elements, except _________________________, have more than one electron that can be removed. Therefore, they have more than one ionization energy. The energy that is needed to remove a second electron is called the ________________________ionization energy. The energy needed to remove a third electron is called the ________________________ionization energy, and so on.
Summarizing Trends for Ionization Energy
Upon examination of the periodic trend for ionization energy it is the opposite of the trend for atomic radius.
Atomic Radius vs First Ionization Energy
Although there are a few exceptions, there are two general trends for ionization energy:
• Ionization energy tends to ________________________down a group as electrons in the outer energy level are further from the positive force of the nucleus, thus they are easier to remove than electrons in lower energy levels.
• Ionization energy tends to ________________________across a period. As you go across a period, the attraction between the nucleus and the electrons in the outer energy level increases. Thus, ______________________ energy is required to pull an electron away from its atom. For this trend to be true, you would expect the noble gases to have the highest ionization energies, and as is shown in Figure 2.16, they do.
Practice Problems
8. Using only a periodic table, rank the elements in each set by increasing ionization energy. Explain your answers. Explanation
(a) Xe He Ar
(b) Sn In Sb
(c) Sr Ca Ba
(d) Kr Br K
(e) K Ca Rb
(f) Kr Br Rb
9. Using only a periodic table, identify the atom in each of the following pairs with the lower first ionization energy.
(a) B O
(b) B In
(c) I F
(d) F N
(e) Ca K
(f) B Tl
Trends for Electron Affinity
Define electron affinity:
Electron affinity is the change in ________________________that occurs when an electron is added to the outer energy level of an atom to form a negative ion.
Figure 2.18 above gives the electron affinities of the main group of elements. If the ion that is formed by gaining an electron is stable, the electron affinity is expressed as a negative integer. The more ________________________ the ion, the higher is the negative integer for the electron affinity. Notice that fluorine has the highest electron affinity. This indicates that fluorine is very likely to be involved in chemical reactions. In fact, fluorine is the most ________________________of all elements. Metals have ___________ electron affinities. That is especially true for the Group 1 (1A) and 2 (2A) elements. Atoms of these elements form ________________________positive ions. A negative ion that is formed by elements of these groups is unstable. It breaks apart into a neutral atom and a free electron Periodic Trend for Electronegativity (p.p. 70 – 71)
Describe. the meaning of electro negativity
What symbol is used to represent electro negativity?
Compare the trend for electroonegativity with that of atomic radius .Summarise the trends.
Explain the trends for electro negativity and atomic radius>
Section Review
1. How does your knowledge of electron arrangement and forces in atoms help you explain the following periodic trends ?
a. atomic radius b. ionization energy c. electron affinity
2. Using only their position on the periodic table, rank each of the following sets of elements in order of increasing atomic size. Explain your answer in each case.
a. Mg S Cl b. Al B In
c. Ne Ar Xe d. Rb Xe Te
e. P Na F f. O S N
3. Using only their position on the periodic table, rank each of the following sets of elements in order of decreasing ionization energy. Explain your answer in each case.
a. Cl Br I b. Ga Ge Se
c. K Ca Kr d. Na Li Cs
e. S Cl Br f. Cl Ar K
4. Which element in each of the following pairs will have the lower electron affinity ? Explain your answer.
a. K or Cl b. O or Li
c. S or Se d. Cs or F
3.1
Classifying Chemical Compounds (p.p. 66 – 74) Most elements do not exist in nature in their pure form as ________________________. Gold, silver, and platinum are three metals that can be found in the earth’s crust as elements. They are called “precious metals” because their occurrence is so rare. Most other metals, and most other elements, are found in nature only as ________________________. There are only about 90 naturally occurring elements. In comparison, there are thousands upon thousands of different compounds in nature, and more are constantly being discovered. Because there are so many compounds, chemists have developed a classification system to organize them according to their properties, such as melting point, boiling point, hardness, conductivity and solubility.
Properties of Ionic and Covalent Compounds Based on their physical properties, compounds can be classified into two groups: ________________________compounds and ________________________compounds. Complete the properties of ionic and covalent compounds in the table below:
PROPERTY
IONIC COMPOUND
COVALENT COMPOOUNDS
State at room temperature
Melting Point
Electrical conductivity as a liquid
Solubility in water
Conducts electricity when dissolved in water
What is Bonding ?
What is a chemical bond?
Which electrons are involved in bonding?
Why do atoms form bonds?
What is responsible for the different properties of ionic and covalent bonds?
Describe the formation of an ionic bond.
Describe the formation of a covalent bond.
Predicting Bond Type Using Electronegativity (p. 72 Text)
How can you use electro negativity to predict bond type?
The Range of Electronegativity Differences
What is a pure covalent bond?
Describe the formation of an ionic bond.
Describe how chemist’s use electronegativity to classify bonds.
What is a polar covalent bond?
Practice Problems
1. Determine the ΔEN for each bond shown. Indicate whether each bond I ionic or covalent
a. O - H
b. C - H
c. Mg - Cl
d. B - F
e. Cr - O
f. C - N
g. Na - I
h. Na - Br
Section Review
1. Name the typical properties of an ionic compound. Give two examples of ionic compounds.
2. Name the typical properties of a covalent compound. Give two examples of covalent compounds.
3. Describe and explain the periodic trend for electronegativity.
4. Based only on their position in the periodic table, arrange the elements in each set in increasing attraction for electrons in a bond.
5. Determine the ΔEN for each bond. Indicate whether the bond is ionic or covalent.
a. N - O b. Mn - O c. H - Cl d. Ca - Cl
3.2 IONIC AND COVALENT BONDINNG: THE OCTET RULE
The Octet Rule When atoms form bonds, they are often more stable. We know that noble gases are the most stable elements in the periodic table. They are extremely non-reactive, and they tend not to form compounds. What the noble gases have in common is a _____________________ outer electron energy level. When an atom loses or gains electrons to achieve a filled outer electron energy level, the atom often becomes more stable. According to the octet rule, atoms bond in order to achieve an electron configuration that is the same as the electron configuration as a noble gas. When two atoms or ions have the same electron configuration, they are said to be isoelectronic. For example, Cl - is isoelctronic with Ar because both have 18 electrons and a filled outer energy level.
Define Isoelectronic:
Ionic Bonding Sodium has a very low electornegativity, and chlorine has a high electronegativty. Therefore, when sodium and chlorine interact, sodium transfers its valence electron to chlorine. As shown in the figure below, sodium becomes Na + 1 and chlorine becomes Cl -. When neutral sodium loses its one valence electron to chlorine, the resulting Na + 1 cation has an en electron energy level that contains ___________ electrons. It is _____________________ with the noble gas neon. On the other hand, chlorine’s outer energy level has seven electrons. When chlorine gains sodium’s electron, it becomes an anion _____________________ with the noble gas argon.
The figure below shows how to represent the formation of an ionic bond using Lewis Structures. Thus, in an ionic bond, electrons are _____________________ from one atom to another so that they form oppositely charged ions. The strong force of attraction between oppositely charged ions is what holds them together.
Practice Problems
1. For each bond below, determine the ΔEN. Is the bond ionic or covalent ? a. Ca - O
b. K - Cl
c. K - F
d. Li - F
e. Li - Br
f. Ba - O
2. Draw Lewis structures to represent the formation of each bond in Question 1
COVALENT BONDING
We have just learned what happens when the electronegativity difference between two atoms is greater than 1.7. The atom with the lower electronegativity transfers its valence electron(s) to the atom with the higher electronegativity. The resulting ions have opposite charges. They are held together by a strong ionic bond. Describe the bond formation between two atoms when the electronegativity difference is zero.
Describe how a coavalent bond is formed.
When two atoms of the same element form a bond, they share their electrons equally in a _____________________ covalent bond. Elements that bond to each other in this way are known as _____________________ molecules. When atoms such as carbon and hydrogen bond to each other, their electronegativities are so close that they share electrons almost equally. Carbon and hydrogen have an electronegativity difference of only 2.6 -= 2.2 = 0.4. In the figure below, you can see how one atom of hydrogen forms a covalent bond with four atoms of hydrogen. The compound methane, CH4 , is formed. Each hydrogen shares one of its electrons with carbon. The carbon shares its four electrons with each hydrogen atom. Thus, each hydrogen achieves a filled outer energy level, as does carbon. (Recall that elements in the first period only need two electrons to fill their outer energy level). When analyzing Lewis structures that show covalent bonds, count the shared electrons as if they belong to each of the bonding atoms.
Practice Problems
1. Show the formation of a covalent bond between two atoms of each diatomic element. a. iodine
b. bromine
c. hydrogen
d. fluorine 2. Use Lewis structures to show the simplest way in which each pair of elements forms a covalent bond, according to the octet rule. a. hydrogen and oxygen
b. chlorine and oxygen
c. carbon and hydrogen
d. iodine and hydrogen
e. nitrogen and hydrogen
f. hydrogen and rubidium
Mutliple Covalent Bonds (p. 82 Text) In covalent bonding, atoms sometimes need to share two or three pairs of electrons, according to the octet rule.
Describe how oxygen gas satisfies the octet rule.
Note: the number of bonds each atom must form is equal to its ________________ Section Review
1. Use Lewis structures to show how each pair of elements forms an ionic bond.
a. magnesium and fluorine b. potassium and bromine
c. rubidium and chlorine d. calcium and oxygen
2. Use Lewis structures to show how each pair of elements forms covalent bonds.
a. one silicon atom and two oxygen atoms
b. one carbon atom, one hydrogen atom, and three chlorine atoms
c. two nitrogen atoms
d. two carbon atoms bonded together with three hydrogen atoms bonded to one carbon atom and one hydrogen atom and one oxygen atom bonded to the other carbon.
3. Use what you know about electronegativity differences to decide what kind of bond would form between each pair of elements.
4. In general, the farther away two elements are form each other on the periodic table, the more likely they are to participate in ionic bonding. Do you agree with this statement? Explain why or why not.
Polar Covalent Bonds: The “In-Between Bonds”
When two atoms have an elecronegativity difference hat is greater than 0.5 but less than 1.7, they are considered to be a polar covalent bond. In a polar covalent bond, the atoms have significantly different electronegativities. The electronegativity difference is not great enough, however, for the less electronegative atom to transfer its valence electrons to the other, more electronegative atom. The difference is great enough for the bonding electron pair to spend more time near the more electronegative atom. For example, the bond between oxygen and hydrogen in water has an electronegativity difference of 1.24. Because this value is less than 1.7 but greater than 0.5, the bond is a polar covalent bond. The oxygen attracts the electrons more strongly than the hydrogen. Therefore, oxygen has a slightly negative charge and the hydrogen has a slightly positive charge. These charges are not full charges but rather a partial charge symbolized by the symbols δ + and δ -.
Practice Problems Predict whether each bond will be covalent, polar covalent, or ionic.
a. C – F
b. O – N
c. Cl – Cl
d. Cu – O
e. Si – H
f. Na – F
g. Fe – O
h. Mn – O
2. For each polar covalent bond in Question 1, indicate the location of the partial charges.
3. Arrange the bonds in each set in order of increasing polarity. (A completely polarized bond is an ionic bond)
a. H – Cl O – O Na – Cl
b. C – Cl Mg – Cl P – O N - N
3.4
Appendix 1 Periodicity and Bonding Questions
Unit 2
CHEMICAL
REACTIONS
CHAPTER 4: CLASSIFYING CHEMICAL REACTIONS: CHEMICALS IN BALANCE
4.1 Chemical Equations Any substance that undergoes a chemical reaction is called a _____________________. A substance that is formed is called a _____________________. Chemists use chemical equations to communicate what is happening in a chemical reaction. Chemical equations come in several forms. Word Equations
A word equation identifies the reactants and products by name. For example, sodium + chlorine → sodium chloride In this equation, “+” means “reacts with” and “→” means “to form”. Practice Problems
1. Describe each reaction using a word equation. Label the reactants and products.
a. calcium and fluorine react to form calcium fluoride
b. barium chloride and hydrogen sulphate react to form hydrogen chloride and barium sulphate
c. calcium carbonate, carbon dioxide and water react to form calcium hydrogen carbonate
d. hydrogen peroxide reacts to form water and oxygen
e. sulfur dioxide and oxygen react to form sulfur trioxide
2. What is one limitation of word equations?
Skeletal Equations Describe a skeletal equation.. Give an example.
A skeletal equation is more useful to a chemist than a word equation because it show the formulas of the compounds involved. It also shows the state of each substance.
Complete the table below that gives the meaning of the symbols in a chemical equation.
SYMBOL
+
→
(s)
(l)
(g)
(aq)
MEANING
Practice Problems
20. Write a skeletal equation for each of the following:
a. solid zinc reacts with chlorine gas to form solid zinc chloride
b. solid calcium and liquid water react to form solid calcium hydroxide and hydrogen gas.
c. Solid barium reacts with solid sulfur to produce solid barium sulfide
d. Aqueous lead (II) nitrate and solid magnesium react to form aqueous magnesium nitrate and solid lead.
21. In each reaction below, a solid reacts with a gas to form a solid. Write a skeletal equation for each reaction.
a. carbon dioxide + oxygen → calcium carbonate
b. aluminum + oxygen → aluminum (III) oxide
c. magnesium + oxygen → magnesium oxide
Why Skeletal Equations Are Incomplete Although skeletal equations are useful, they do not fully describe chemical reactions. For example, according to the skeletal equation showing the formation of sodium chloride, one molecule of sodium reacts with one chlorine molecule containing two chlorine atoms. The product is one formula unit of sodium chloride containing one atom of sodium and one atom of chlorine. Where has the extra chlorine atom gone?
State the Law of Conservation of Mass
Balanced Chemical Equations A balanced chemical equation reflects the Law of Conservation of Mass. This type of equation shows that there is the same number of each type of atom on both sides of the equation. Some skeletal equations are already balanced. For example, the skeletal equation for the reaction between carbon and oxygen to form carbon dioxide shows one carbon and two oxygen atoms on the reactant side of the equation, and one carbon and two oxygen atoms on the product side of the equation
Most chemical equations, however, are not balanced, such as the one showing the formation of sodium chloride. Na + Cl2 → NaCl
You cannot balance an equation by changing any of the chemical formulas. The only way to balance an equation is to put the appropriate numerical coefficient in front of each compound or element in the equation.
Practice Problems
22. Copy the following skeleton equations and balance them.
a. S (s) + O 2 (g) → SO 2 (g)
b. P 4 (s) + O 2 (g) → P 4 O 10 (s)
c. H2 (g) + Cl 2 (g) → HCl (g)
d. SO 2 (g) + H2 O (l) → H2 SO4 (aq)
23. Indicate whether these equations are balanced. If they are not, balance them.
a. 4 Fe(s) + 3 O 2 (g) → 2Fe2O3(s)
b. HgO(s) → Hg (l) + O 2 (g)
c. H2 O2 → 2 H2 O (l) + O 2 (g)
d. 2 HCl (aq) + Na2SO3 (aq) → 2 NaCl (aq) + H2 O (l) + SO 2 (g)
NOTES
Steps for Balancing Equations Here are some steps to follow when balancing equations that are more complex:
1. Balance the element, other than hydrogen and oxygen that has the greatest number of atoms in reactant or product.
2. Balance the other elements other than hydrogen and oxygen
3. Balance hydrogen and oxygen, whichever is present in the combined state. Leave until last, whichever one is in uncombined state
4. Check that each equation is balanced by counting the number of atoms of each element on each side of the equation
When the equation is balanced, the coefficients should be in their lowest terms.
Practice Problems (text p.p. 117 & 118) # 7, 8, 9
Section Review (text p. 118) # 2, 4
Synthesis and Decomposition Reactions
(text p.p. 119 – 125)
Just as there are different types of compounds are different types of chemical reactions. We will learn about five major classifications used for chemical reactions.
Synthesis Reactions
In a synthesis reaction, two or more elements or compounds combine to form a new substance.
Write the general equation for a synthesis reaction
In this course all you'll the with two specific types synthesis reactions involving compounds that you are familiar with -- oxides and water.
Describe the products formed when a non-metallic oxide reacts with water.
Describe the acids that are formed when a nonmetallic oxide and water react. Provide an example.
Describe the products formed when a metallic oxide reacts with water. Provide an example and write the balanced equation.
Practice Problems p. 122 # 10 – 13
Decomposition Reactions
Write the general formula for a decomposition reaction.
C → A + B
What is a decomposition reaction? Provide an example and a balanced equation
Practice Problems p.. 123 # 14 - 16
Combustion Reactions
Describe the products of a complete combustion reaction. Provide an example.
What is incomplete combustion?
Practice Problems p. 124 # 17- 20
Section Review P. 125 # 1 – 6
4.3
Single Displacement and Double Displacement Reactions
Describe a single displacement reaction.
Write the general reactions that represent two different types of single displacement reactions. Describe these two reactions.
Single Displacement Reactions and the Metal Activity Series
Most single displacement reactions involve one metal displacing another metal from a compound. In the following equation, magnesium metal replaces the zinc in ZnCl2, thereby liberating the zinc as the metal.
Mg(s) ZnCl2 (aq) → MgCl2 (aq) + Zn(s)
Writes three balanced equations that illustrate the various types of single displacement reactions involving metals. 1.
2.
3.
Write the general formula for a single displacement reaction
Write three guidelines that you should follow when analyzing single displacement reactions
•
•
•
Practice problems # 21 (text p. 127)
What is at activity series
Through experimentation chemists have ranked the we reactivity of the metals including hydrogen [in acids and water], in ant activity series.
Complete Investigation 4-A Creating an Activity Series of Metals
p.p. 128-129
The Metal Activity Series
Describe the arrangement of metals on the metal activity series
What does a single displacement reaction always favor?
Using the activity series predict the products of the following reactions:
Fe(s) + CuSO4 (aq) → FeSO4 (aq) + Cu(s)
Ag (s) + CaCl2(aq) →
Practice Problems (text p. 131 #22)
Single Displacement Reactions Involving a Halogens
Predict to the products for the following reaction:
Cl2 (g) + 2 KBr (aq) →
Describe the activity series for halogens.
Predict the products of the following reactions:
F2(g) + 2NaCl →
I2(g) + CaBr2(aq) →
Practice Problems: (text p. 131 #23, 24)
Single Displacement Reactions Involving Halogens Non-metals, typically halogens, can also take part in single displacement reactions. For example, molecular chlorine can replace bromine from KBr, an ionic compound, producing bromine and potassium chloride.
Cl2(g) + 2 KBr(aq) → 2 KCl(aq) + Br2(l)
The activity series for the halogens directly ________________________the position of the halogens in the periodic table. It can also be shown in the following way. ________________________ is the most reactive, and ________________________is the least reactive.
F > Cl > Br > I
The activity series for halogen can be used the same way the activity series for metals is used. If the element is above another in the activity series, it will replace it. For example,
F2 (g) + 2NaCl (aq) → 2NaF (aq) + Cl2 (aq)
If the element is below another in the activity series, it will not replace it and no reaction will occur. For example,
I2 + CaBr2 (aq) → no reaction
Practice Problems (Text p. 131)
Double Displacement Reactions
Describe what takes place in a double displacement reaction.
Write the general formula for a double displacement reaction.
An example of a double displacement reaction between sodium chloride and silver nitrate
Since silver chloride is virtually insoluble in water, it forms a solid compound or precipitate.
Double displacement reactions tend to occur in aqueous solutions. You can tell a double displacement reaction has taken place in the following cases:
What are three pieces of evidence that a double displacement reaction has occurred?
•
• •
Double Displacement Reactions that Form a Precipitate
What is a precipitate ?
How can you predict whether a precipitate will be produced during a reaction?
Given the following reactants, how would you predict the products of the reaction and their state ? (Note, many hydroxide compounds are insoluble. Potassium cations form soluble substances with all anions)
MgCl2(aq) + KOH(aq) →
If both products are soluble, and neither product precipitates out, both ionic compounds are dissolved in water and no reaction occurs.
Practice Problems (text p. 134)
Double Displacement Reactions that Produce a Gas In certain cases, you know that a double displacement reaction has occurred because a gas is produced. The gas is formed when one of the products of the double displacement reaction decomposes to give water and a gas. For example, the reaction between sodium carbonate and hydrochloric acid. If you carry out this reaction, you immediately see the formation of carbon dioxide gas. The first reaction that takes place is a double displacement reaction. Determine the products in the following way.
• Separate the reaction into ions, and switch the anions. Write the chemical formulas for the products and balance the equation.
Na2CO3 (aq) + 2HCl (aq) → 2NaCl (aq) + H2CO3 (aq)
The carbonic acid is unstable and decomposes to carbon dioxide and water.
H2CO3 (aq) → H2O (l) + CO2 (g
Practice Problems (text p. 135 )
The Formation of Water In Neutralization
Neutralization reactions are a special type of double displacement reaction that produces water.
What is a neutralization reaction ? Give an example.
Often, neutralization reactions produce no precipitate or gas.
Practice Problems (text p. 135)
Investigation 4-B (text p.p. 136 – 137)
Investigation 4-c (text p. 138 )
Section Review Questions (text p. 140
Chapter 4 Review text p. 149
Unit 1 review (text p. 156)
Date: Name: Class:
STSE Connections in the Great Lakes
Goal
Outline the connections among science, technology, society, and the environment in relation to the Great Lakes.
Procedure
1. Choose an Area of Concern or a more general issue from those listed below. Use print and electronic resources to investigate the chemicals involved, and the connections between science, technology, society, and the environment.
Great Lakes Areas of Concern
• Hazardous waste sites along the St. Lawrence river in Cornwall, Ontario, and in the Port Hope harbour, remain from past industrial waste handling practices.
• Sewage treatment plants and sewage overflow contributes to the pollution in the Bay of Quinte.
• Copper mining, milling, and smelting operations have contaminated the water in Torch Lake, connected to Lake Superior.
• Industrial chemical discharges pollute the Niagara River connecting Lake Erie and Lake Ontario.
• Urban run-off from Metro Toronto contaminates the water in Lake Ontario.
General Pollution Issues involving the Great Lakes
• The atmospheric/airborne deposition of toxic chemicals may have a large role in the pollution of the Great Lakes.
• Biomagnification and bioaccumulation of toxic chemicals through the food chain puts predators such as lake trout, salmon, herring gulls, and fish-eating humans at risk. (Hint: Check out fish consumption advisories for the Great Lakes.)
• Nuclear power plants discharge heated water into the Great Lakes. How does this affect the environment? Where is this happening?
2. Organize and present your findings using one or more media such as an essay, a web site, a PowerpointTM or CorelTM presentation, or a brochure.
Analysis
1. What chemicals are involved in the issue you investigated?
2. For what purpose were the chemicals originally produced?
3. How did the chemicals benefit society?
4. How did the chemicals affect the environment?
5. What is being done to solve the problem?
Conclusions
6. Use a concept web to illustrate the STSE connections present in this issue.
UNIT C
QUANTITIES
IN
CHEMICAL
REACTIONS
5.1
Isotopes and Average Atomic Mass The mass of an atom is expressed in atomic mass units (u.). Atomic mass units are a relative measure, define by the mass of one carbon – 12 atom. One carbon – 12 atom has been assigned a mass of 12 u, or 1u = 1/12 the mass of one atom of carbon – 12. Average Atomic Mass
What is the average atomic mass of an element?
6 C 12.01
Atomic Weights
Provide the formula to calculate atomic weight. What do the symbols represent?
Sample Problem
Naturally occurring silver exists as two isotopes. From the mass of each isotope and the isotopic abundance listed below, calculate the average atomic mass of silver. ISOTOPE ATOMIC MASS (U) RELATIVE ABUNDANCE Ag-107 106.9 51.8 Ag – 109 108.9 48.2
Solution
Practice Problems p.167 & 168 # 1- 4
5.2
THE AVOGADRO CONSTANT AND THE MOLE
(p. 171 Text)
Mole (symbol mol). What number does the mole represent?
•
•
How many formulae units would one mole NaCl contain?
Practice Problems
11. If you drove for 6.022 x 10 23 days at a speed of 100 km/h, how far would you travel ?
12. If you spent $ 6.022 x 10 23 at a rate of $1.00 /s, how long, in years would the money last? Assume that every year has 365 days.
Converting Moles to Number of Particles Write the Formula Converting Moles to Number of Particle
Example Problem
A sample contains 1.25 mol of nitrogen dioxide NO2.
a. How many molecules are in the sample ?
b. How many atoms are in the sample ?
Solution
Practice Problems (text p. 177 )
Converting Number of Particles to Moles Write the Formula Converting Number of Particles to Moles
Sample Problem: How many moles are present in a sample of carbon dioxide, CO2, made up of 5.83 x 10 24 molecules? Solution
Practice Problems (text p. 178)
MOLES WORKSHEET
1. What is meant by the term of mole ?
2. What is meant by the term Molar Volume?
3. What Is meant by the term of Molar Mass?
4. How many atoms make up one molecule of the following?
a. NaCl b. CCl4 c. H3PO4
d. Fe2(SO4)3 e. Al2(Cr2O7) 3 f. Ca(OH) 2
5. Determine the Molar Mass each of the following
a. H2SO3 b. C3H8 c. Ca(HCO3) 2
d. MnSO4 e. U(CH3COO) 3
6.
a. How many molecules are present and 5.2 moles of H2O
b. How many atoms are present and 5.2 moles of H2O
c. How many molecules are present in three moles of NH3
d. How many activities are present in three moles of NH3
5.3
Molar Mass
One mole of an element has a mass expressed in grams numerically equivalent to the element’s average atomic mass expressed in atomic mass units. One mole of zinc atoms has a mass of 65.39 g. You can use the periodic table to determine the mass of one mole of an element. What is Molar Mass
Complete the table below
Element
Average Atomic Mass (u)
Molar Mass (g)
Hydrogen, H
Oxygen, O
Sodium, Na
Argon, Ar
The molar mass relates the amount of an element or compound in moles, to its mass.
Finding the Molar Mass of Compounds
Find the molar mass of BaO
Sample Problem
What is the mass of one mole of calcium phosphate, Ca3(PO4)2 ? Solution
Practice Problems p. 184 Text # 23 - 26
Investigation 5-A (text p.p. 182 & 183)
Complete Investigation 5-A and answer Analysis 1-6, Conclusion 7, and
Application 8, 9 in this workbook.
Converting from Moles to Mass
Write the formula that converts moles to mass.
Sample Problem
A flask contains 0.75 mol of carbon dioxide gas, CO2. What mass of carbon dioxide gas is in the sample ?
Solution
Practice Problems (text p. 186 text)
The triangle shown below is useful for problems involving number of moles, number of particles, and molar mass.
Converting from Mass to Moles
Write the formula used to convert mass to moles
Sample Problem
How many moles of acetic acid, C2H3COOH, are in a 23.6 g sample ? Solution
Practice Problems p. 187 text # 31 – 34
Sample Problem - Particles to Mass
What is the mass of 5.67 x 1024 molecules of cobalt (II) chloride, CoCL2
Solution
Practice Problems (text p. 190 # 35 – 38)
Sample Problem - Mass to Particles
Chlorine gas, Cl2, can react with iodine, I2 to form iodine chloride, ICl. How many molecules of iodine chloride are contained in a 2.74 x 10 –1g sample ? Solution
Practice Problems p. 191 text # 39 - 42
Ch. 5 Review p. 193 text
Mole Calculations Worksheet
1. Calculate the number of moles of the stated as substance of each of the following a samples
a 225 g of aspirin, C 9H8O4
b. 35 g of baking soda NaHCO3
c. 1.45 kG of gold
d. 0.84 g of hydrogen gas
e. 163 g sodium fluoride
f. 46.7 g of methanol, CH3OH
2. Calculate the mass of the following:
a. 12.4 mol of helium gas, He
b. 0.26 mol of butane. C4H10
c. 255 mol of ammonia
d. 1.8 mol of magnesium hydroxide
Answers
1a. 1.25 mol b. 0.417 mol c. 7.36 mool d. 0.42 mol
e. 3.88 mol f. 1.46 mol
2a. 49.6 g b. 15 g c. 4340 g d. 100 g
PECENTAGE COMPOSITION
(text p. 199 – 206)
State the law of definite proportions
carbon monoxide, CO, carbon dioxide, CO2,
C - 42.88 % by mass C – 27.29% by mass The percentage composition of a compound refers to the relative mass of each element in the compound.
Sample Problem
A compound with a mass of 48.72 g is found to contain 32.69 g of zinc and 16.03 g of sulfur. What is the percentage composition of the compound ?
• Mass percent of zinc =
• Mass percent of S =
Practice Problems . (text p. 201)
Calculating Percentage Composition from a Chemical Formula (p. 202 – 204) Percentage composition can be calculated from a known chemical formula. If you assume you have one mole of compound, you can use the molar mass of the compound, with its chemical formula, to calculate its percentage composition.
Example: Find the percentage composition by mass of HgS
Solution:
Sample Problem: Cinnamaldehyde, C9H8O, is responsible for the characteristic odour of cinnamon. Determine the percnetage composition of cinnamaldehyde by calculating the precents of carbon, hydrogen, and oxygen.
Solution:
Practice Problems (text p. 204 # 5 – 8)
Section Review p. 205 & 206
PERCENT COMPOSITION
WORKSHEET
Part A CALCUATING PERCENTS
When elements combine to form a given compound they are always present with same composition by mass
KClO3
% K % Cl % O
Fe(OH) 3
% Fe % O % H
Part B Calculating Simplest Formula (Empirical formula)
When elements combine to form compounds, they do so simple whole numbers.
Sample Problem
11.5 grams of sodium burdens in air to form 15.5 g of sodium oxide. Calculate the simplest formula of sodium oxide.
Na O mass # moles
Simplest mole ratio simplest formula Sample Problem
A compound consists of 27.29% carbon and 72.71% oxygen by mass. Calculate the simplest formula of this compound
C O
%
Mass in 100 g
# moles
Simplest mole
Ratio
Simplest formula Sample Problem
A compound consists of 52.17% carbon, 13.04% hydrogen and 34.78% oxygen. Calculate the simplest formula.
What does the empirical formula show?
What is the molecular formula of a compound?
Complete the following table:
Name of Compound
Molecular Formula
Empirical Formula
Lowest Ratio of Elements
Hydrogen peroxide
Glucose
Benzene
Acetylene (ethyne)
Aniline
Water
As you can see, it is possible for different compound to have the same empirical formula.
Determining a Compounds Empirical Formula (p. 208)
Sample Problem: Calculate the empirical formula of a compound that is 85.6 % carbon and 14.4% hydrogen.
Solution
Practice Problems p. 209 # 9 -15
Investigation 4-A (text p.p.212. 213)
Practice Problems (text p. 218)
7.1
STOICHIOMETRY
(text p.p. 234 – 259)
Balanced chemical equations are essential for making calculations related to chemical reactions.
Earlier, we learned how to classify types of chemical reactions and balanced chemical equations that describe them. We have also learned how to relate the number of particles and a substance to the amount of a substance in moles and grams.
Express Lab (text p. 235)
Complete the express Lab, Mole Relationships in a Chemical Reaction,
Answer Analysis 1 – 6
Particle Relationships in a Balanced Chemical Equation (text p.p. 235 – 237)
Recall that the coefficients in front of compounds and elements chemical equation tell us how many atoms and molecules participate in any reaction. For example, in the Haber
Process, where ammonia is prepared industrially from its elements:
N2(g) + 3H2(g) → 2NH3 (g)
This the equation tells us that one molecule of nitrogen gas reacts with three molecules of hydrogen gas to form two molecules of ammonia gas. You can use a ratio to express the number of atoms in the equation:
1 molecule N2 (g): 3 molecules H2(g) 2 molecules NH3 (g)
Example` If you wanted to 20 molecules of ammonia, how many molecules of nitrogen would you need?
Practice problems [text p. 237]
Mole Relationships in Chemical Equations
Until now, we have assumed that the coefficients in a chemical equation represent particles.
They can, however, represent moles. For example, from the Haber process shown earlier, shows that 1 mole of nitrogen reacts with 3 moles of hydrogen to form 2 moles of ammonia.
You can manipulate the mole ratios in the same way you manipulate ratios involving molecules. Sample Problem Using the equation above, how many moles are produced by 2.8 moles of hydrogen?
Practice problems [text p. 238)
Different Ratios of Reactants (text p. 239)
The relative amounts of reactant is are important. For example, carbon can combine with oxygen in 2 different ratios forming carbon dioxide, CO2, a product of cellular respiration in animals and one of the products of the complete combustion of a hydrocarbon fuel or
CO, a colourless, tasteless, odourless, highly poisonous gas that is responsible for hundreds of deaths in Canada and the United States each year.
Example Vanadium to form several different compounds with oxygen including V2O5, VO2, and V2O3. Determining the number of moles of oxygen that are needed to react with 0.56 mol of vanadium to form divanadium pentoxide.
Practice Problems [text p. 240]
Mass Relationships in Chemical Equations (text p. 241 – 246)
You have just learned that the coefficients in balanced equations represent moles as well as particles. Therefore, you can use the molar masses of reactants and products to determine the mass ratios for a reaction. Stoichiometric Mass Calculations
If you know what the amount of one substance in a chemical reaction [in particles, mass, or moles], you can calculate the amount of any other substance in the reaction to [in particles, mass, or moles], using the information in the balanced chemical equation.
Balanced chemical equations allow chemists to predict the amount of products that will result in a reaction involving a known amount of reactants. As well, chemists can use a balanced equation to calculate the amount of reactants they will need to produce the desired amount of products. They can also use it to predict the amount of one reactant they will need to completely react with another reactant.
Sample Problem: MASS TO MASS CALCULATIONS FOR PRODUCTS AND REACTANTS
Problem: Hydrazine, N2H4, and dinitrogen tetroxide, N204, formed a fuel mixture used to launch a lunar module. These two compounds react to form nitrogen gas and water vapour. If 50.0 g of hydrazine reacts with sufficient dinitrogen tetroxide, what mass of hydrogen gas is formed ?
Solution
i. Write a balanced chemical equation
ii. Convert mass of hydrazine to number of moles of hydrazine
iii. Calculate the number of moles of nitrogen, using the mole ratio of hydrazine to nitrogen iv. Convert the number of moles of nitrogen to grams
Practice Problems (text p. 246)
A GENERAL PROCESS FOR SOLVING STOICHIOMETRIC PROBLEMS
Step 1 Write a balanced chemical equation
Step 2 If you are given the mass or the number of particles, convert it to the number of moles
Step 3 Calculate the number of moles of the required substance based on the number of moles of the given substance, using the appropriate mole ratio.
Step 4 Convert the number of moles of the required substance to mass or number of particles as directed by the question
Sample Problem MASS AND PARTICLE STOICHIOMETRY
Problem: Passing chlorine gas through molten sulfur produces liquid sulfur dichloride. How many molecules of chlorine react to produce 50.0 g of sulfur dichloride ?
Solution
i. Write a balanced chemical equation
ii. Convert the given mass of sulfur dichloride to the number of moles
iii. Calculate the number of moles of chlorine gas using the mole ratio of chlorine to sulfur dichloride
iv. Convert the number of moles of chlorine gas to the number of particles of chlorine gas.
Practice Problems (text p. 248 bottom & 249 top)
Stoichiometry Problems Types 1-5
N2(g) + 3H2(g) → 2NH3 (g)
1. Mole -- Mole
Calculate the number of moles of ammonia [NH3 (g)] that can be produced from 5.6 moles of hydrogen gas [H2(g) ].
2. Mole – Mass
Calculate the number of grams of nitrogen gas (N2(g)) needed to react with five moles of hydrogen gas (H2(g))..
3. Mole – Volume (STP)
Calculate the number of moles of ammonia that would be produced from 8 litres of hydrogen gas STP
4. Mass – Mass Problems
a. Determine the mass of copper that can be produced from 217 g of copper (II) chloride
CuCl2 → Cu + Cl2
b. Iron reacts with oxygen gas to form iron (III) oxide. Determine the mass of iron (III) oxide produced from 112 grams of iron, given:
4 Fe + 3 O2 → 2Fe2O3
c. Calculate the mass of oxygen gas produced when 43.8 g of potassium chlorate is decomposed by heating, given:
2 KClO3 → 2 KCl + 3 O2
d. Calculate the mass of the BaSO4 produced from 24.1 grams of BaCl2 given:
BaCl2 + Na2SO4 → BaSO4 of + 2 NaCl
5. Mass – Volume Problems
a. Determine not the volume all of nitrogen gas at STP required to produce 12.6 g of ammonia, given N2 + 3H2 → 2 NH3
b. Determine the mass of aluminum required to produce 1.8 Litres of hydrogen gas at STP, given:
2 Al + 6HCl (aq) → 2AlCL3 + 3 H2
c. Magnesium reacts with HCl to produce hydrogen gas and magnesium chloride. Calculate the mass of magnesium required to produce 15 litres of hydrogen gas at STP , given:
Mg + 2 HCl → MgCl2 + H2
STOICHIOMETRY : TYPES 1-5
1. Fe + 2 HCl -------> FeCl2 + H2
(a) How many moles of iron would react with 6 moles of HCl ?
(b) How many moles of HCl would react with 0.1 moles of Fe ?
(c) How many grams of iron would produce 2 moles of H2?
(d) How many grams of iron would produce 300 grams of H2?
(e) How many grams of iron would produce 700 g of FeCl2?
(f) How many grams of iron would produce 22.4 L of H2 at STP ?
(g) How many grams of iron would produce 300 L of H2 at STP ?
2. C3H8 + 5 O2 ---------> 3 CO2 + 4 H2O
(a) How many moles of oxygen are required if 6 moles of C3H8 is to totally react?
(b) How many moles of CO2 will be produced is 15 moles of oxygen totally reacts?
(c) How many moles of water are formed if 12.2 moles of oxygen totally react ?
(d) What mass of oxygen will react with 2.5 moles of C3H8 ?
(e) What mass of carbon dioxide is formed if 30 grams of propane is totally reacted?
(f) What mass of carbon dioxide is formed if 263 grams of water is formed ?
(g) What mass of water is formed if 7800 grams of propane totally reacts ?
3. 3 H2SO4 + 2 Al ----------> Al2(SO4)3 + 3 H2
(a) How many moles of hydrogen gas are produced if 36.5 moles of sulphuric acid totally reacts?
(b) What mass of aluminum metal would produce 500 Litres of hydrogen gas at STP ?
(c) What mass of H2SO4 would produce 500 L of H2 at STP ?
(d) What mass of H2SO4 would produce 500 g of Al2(SO4)3 ?
(e) What mass of Al2(SO4)3 would be produced if 5786 L of H2 were produced at STP ?
7.2
THE LIMITING REAGENT
(p.p. 251 – 259 text) What are stoichiometric amounts?
In practice, however, there are often reactants left. In nature, reactions almost never have reactants in stoichiometric amounts. For example when an animal carries out respiration,
C6H12O6 (s) + 6O2 (g) → 6 CO6 (g) + 6 H2O (l)
there is an unlimited supply of oxygen in the air. The amount of glucose, C6H12O6 , depends on how much food the animal has eaten.
Determining the Limiting Reactant
What is meant by the limiting reactant?
Sample Problem IDENTIFYING THE LIMITING REACTANT
Problem Lithium nitride reacts with water to form ammonia and lithium hydroxide, according to the following equation:
Li3N (s) + 3H20 (l) → NH3 (g) + 3 LiOH (aq)
If 4.87 g of lithium nitride reacts with 5.80 g of water, find the limiting reagent.
Solution
i. Convert given masses into moles
ii. Use mole ratios of reactants and products to determine how much ammonia is produced by each amount of reactant.
•
• Amount of ammonia produced based on amount of H20
Alternative Method for Determining Limiting Reactant
Practice Problems (text p. 254 )
Inv 7-A (text p. 255)
The Limiting Reactant in Stoichiometric Problems
Since chemical reactions usually occur with one or more of the reactants in excess, it is often necessary to determine the limiting reactant before you carry out stoichiometric calculations.
Sample Problem THE LIMITNG REACTANT IN A STOICHIOMETRIC PROBLEM Problem White phosphorous consists of a molecule made up of four phosphorus atoms. It burns in pure oxygen to produce tetraphosphorus decaoxide.
P4 (s) + 5O2 (g) → P4O10 (s)
A 1.00 g piece of phosphorus is burned in a flask filled with 2.60 x 10 23 molecules of oxygen gas. What mass of tetraphosphorus decaoxide is produced?
Solution
i. Convert each reactant to moles and find the limiting reactant.
ii. Determine the limiting reactant
• Determine the amount of P4O10 that would be produced by 8.07 x 10 - 3 mol P4
• Determine the amount of P4O10 that would be produced by 0.432 mol O2
iv Using the mole to mole ratio of the limiting reactant to the product, determine the number of moles of phosphorus decaoxide that is expected
v. Convert this number of moles to grams
Or alternative solution
Practice Problems (text p.p. 257 & 258)
Section Review (text p.p. 258 & 259)
7.3
PERCENTAGE YIELD
(p.p. 260 – 270 text)
Theoretical Yield and Actual Yield
What does percent yield mean?
•
•
Write the formula used for calculating Percentage Yield
Sample Problem CALCUALTING PERCENTAGE YIELD
Problem: Ammonia can be produced by reacting nitrogen gas, taken from the atmosphere, with hydrogen gas.
N2 (g) + 3H2 (g) → 2NH3 (g)
When 7/5 a 10 1g of nitrogen reacts with sufficient hydrogen, the theoretical yield of ammonia is 9.10 g. If 1.72 g of ammonia is obtained by experiment, what is the percentage yield of the reaction ?
Solution
i. Divide the actual yield by the theoretical yield and multiply by 100 %.
Practice Problems (text p. 262)
Sample Problem PREDICTING ACTUAL YIELD BASED ON PERCENTAGE YIELD
Problem Calcium carbonate can be thermally decomposed to calcium oxide and carbon dioxide.
CaCO3(s) → CaO (s) + CO2 (g)
Under certain conditions, this reaction proceeds with a 92.4 % yield of ca;cium oxide. How many grams of calcium oxide can the chemist expect to obtain if 12.4 g of calcium carbonate is heated ?
Solution
i. Convert from grams of calcium carbonate to moles of calcium carbonate
ii. Use stoichiometry to calculate the theoretical yield of calcium oxide
iii. Convert from mol CaO to grams of CaO
iv. Calculate actual yield by multiplying theoretical yield by percentage yield
Or alternative method
‘
Practice Problems (text p. 264)
Inv 7-B (text p.266)
Complete Investigation 7-B, (text p. p. 255 – 267) and answer Analysis 1, 2, and
Conclusion 3
STOICHOMETRY - TYPES 6-8
6. VOLUME-VOLUME (at the same conditions)
What volume of N2 at 20oC and 730 mm Hg would produce 400 Litres of NH3 at the same conditions?
N2 + 3 H2 ------> 2 NH3
7. VOLUME - MASS (not STP)
What mass of Al is required to form 1800 Litres of H2 at 30oC and 80 kPa ?
2 Al + 3 H2SO4 ------> Al2(SO4)3 + 3 H2
8. MASS- VOLUME (not STP)
What volume of hydrogen gas is formed at 80oC and 0.5 atm if 130 grams of Al metal is reacted?
2 Al + 3 H2SO4 ------> Al2(SO4)3 + 3 H2
SCH3U STOICHIOMETRY: TYPES 6-8
1. 4 Li + O2 ---------> 2 Li2O
What mass of Li would be needed to react with 35 L of oxygen gas at 25 degrees C and 900 mm Hg ?
2. Solid ammonium dichromate, when heated, yields gaseous nitrogen, chromium (III) oxide and water.
(NH4)2Cr2O7 ---------> 4 N2 + Cr2O3 + 4 H2O
What volume of nitrogen gas at 80 degrees C and 104.6 kPa would be formed if 8 grams of water is also formed ?
3. 2 BBr3 + 3 H2 --------> 2 B + 6 HBr
(a) What volume of boron bromide at 300 degrees C and 65 kPa would produce 600 L of hydrogen bromide at the same conditions ?
(b) What mass of boron is formed if 3700 L of hydrogen gas at 1500 degrees C and 1.8 atm is totally reacted ?
4. 2 C10H20O + 29 O2 --------> 20 CO2 + 20 H2O
(a) What volume of oxygen gas is required at 300 degrees C and 800 mm Hg to produce 2500 grams of carbon dioxide ?
(b) What volume of water vapour is produced at 300 degrees C and 85 kPa if 1800 grams of oxygen gas totally reacts ?
(c) What volume of carbon dioxide at 400 degrees K and 1.75 atm will be produced if 2500 grams of C10H20O totally reacts?
Name __________________________
A Selection of Enjoyable Stoichiometry Problems
Consider the combustion of ethyl alcohol (C2H5OH) being investigated by a scientist called Barbara.
C2H5OH (l) + 3O2 (g) → 2CO2 (g) + 3H2O(g)
1. How many moles of carbon dioxide are produced when Barbara burns 1.50 moles of C2H5OH and
2. What mass of carbon dioxide is produced 1.5 moles ethanol is burned?
3. What mass of carbon dioxide is produced 23.0 in g of ethanol is burned?
Consider the following reaction for questions 4, 5 and 6
3Ag(s) + 4HNO3(aq) → 3AgNO3 (aq) + NO(g) + 2H2O(l)
4. How many moles of NO are produced when 1.5 moles of Ag reacts with excess HNO3?
5. How many grams of NO are produced when 1.5 moles of Ag reacts with excess HNO3(aq)?
6. How many grams of AgNO3 are produced when 189 g of HNO3 reacts with excess Ag?
Consider the following reaction:
3Cu(s) + 8HNO3 (aq) → 3Cu(NO3) 2(aq) + 2NO(g) + 4H2O(l)
7. How many moles NO are produced the reaction 6.35 g of Cu with excess nitric acid?
8. What mass of water is produced 12.6 g of nitric acid reacts with excess Cu?
MATTER, CHEMICAL TRENDS AND CHEMICAL BONDING
UNIT 1
MATTER AND CHEMICAL BONDING
1.1 The Study of Chemistry
Chemistry is the study of matter and its properties, the changes (chemical, physical, and nuclear) that matter experiences, and the energy that is associated with these changes.
The term STSE refers to science, technology, society, and the environment. These four areas are strongly interconnected. Science and technology, working on behalf of society, have produced chemicals that help every part of our lives. However, something that benefits society may harm the environment and end up harming society by endangering human health in some way. Science and technology working together have created a better quality of life for our society. At the same time, they have caused many environmental problems. Today, science and technology must work together to solve these and other problems.
1.2
Physical properties—colour, density, boiling point, melting point—are properties of matter that are observable without changing one type of matter into another type of matter. Chemical properties—reactivity (with oxygen, with acid, or with other substances), toxicity—are properties that are observable when one type of matter changes into a different kind of matter of different properties.
All measurements contain some degree of uncertainty because the last digit is estimated. The certainty of a calculated result depends on the certainty of the data, which in turn depends on the certainty (degree of accuracy) of the measuring instrument.
Calculated answers have the same number of significant digits as the least-certain measurement involved in the calculation; digits in excess of the number of least-certain digits must be rounded off when expressing the final answer. Precision refers to how close values are to each other, while accuracy refers to how close values are to an actual value.
Generally, the more precise a measurement is, the more significant digits it has.
Investigation 1-A Observing Aluminum Foil (text p. 13)
Complete Investigation 1-A and answer Analysis # 1, 2 Conclusion 3 a, b, c,
Using Measurement to Describe Matter
What is the name of the system of measurement that scientists rely on to communicate effectively. What does it allow scientists to do?
The system of measurement used by scientists is called the International System of Units or (S.I.). It allows scientists anywhere in the world to describe matter in the same quantitative language.
Name five of the base S.I. units?
Five of the base S.I. units are the metre (m), the kilogram (kg), the second (s), the mole (mol and the Kelvin (K).
Complete the following table on Important S.I. Quantities and Their Units Quantity Definition S.I. Units or their derived equivalents Equipment used to measure the quantity
mass
The amount of matter in an object kilogram (kg) gram (g) milligram (mg)
Balance or scale
length The distance between two points metre (m) centimeter (cm) millimetre (mm) ruler temperature The average kinetic energy of a substance
Kelvin )K) degrees Celsius(° C)
Thermometer
volume The amount of space an object occupies cubic metre (m3) cubic centimeter (cm3) litre (L) milliliter (mL)
beaker, graduated cylinder, pipette
Mole the amount of 6.023 x 1023 molecules of a substance mole (mol) calculated, not measured
Density the mass per unit volume of a substance kilograms per cubic metre (kg/m3) grams per cubic centimeter (g/cm3) calculated or measured
Energy the capacity to do work (to move matter)
Joule (J) calculated not measured
Measurement and Uncertainty
What are exact numbers?
Exact numbers are numbers you can count or are true by definition. They are numbers that you are certain about.
What two factors affect your ability to communicate measurements and calculations.
Two main factors affect your ability to communicate measurement and calculations. One factor is the instrument you use, and the other factor is your ability to read and interpret what the instrument is telling you.
Significant Digits, Certainty and Measurement
All measurements involve uncertainty. One source of this uncertainty is the measuring device itself. Another source is your ability to perceive and interpret a reading. The last digit (farthest right) in any measurement is always an estimate.
The digits that you record when you measure something are called significant digits. Significant digits include all the digits that you are certain about and a final, uncertain digit that you estimate.
How Can You Tell which Digits are Significant ?
Complete the following table of rules for determining significant digits:
Rules Examples
1. All non-zero numbers are significant 7.886 has four significant digit.
19.4 has three significant digits
527.266 992 \has nine significant digits
2. All zeros that are located between two non-zero numbers are significant 408 has three significant digits
25 047 has five significant digits
3. Zeros that are located to the left of a value are not significant 0.0907 has three significant digits. They are 9, the third 0 to the right, and the 7. The function of the 0.0 at the beginning is only to locate the decimal.
0.000 000 000 06 has only one significant digit
4. Zeros located to the right of a value may or may not be significant
22 700 may have three significant digits, or it may have five significant digits.
Explaining Three Significant Digits in the Example Above
The Great Lakes contains 22 700 km3 of water, but this an approximate value. Using scientific notation, you can rewrite 22 700 km3 as 2.27 x 104 km3 . This shows that only three digits are significant.
Explaining Five Significant Digits in the Example Above
If you were able to measure out the volume of water in the Great Lakes and verify the value of 22 700 km3 , then all five digits, including the zeros, would be significant. Once again, scientific notation lets you show clearly the five significant digits: 2.2700 x 104 km3
Complete Practice Problems #1, 2 (text p. 18)
Calculating With Significant Digits (text p. 20 – 22)
In this course, you will often take measurements and use them to calculate other quantities. You must be careful to keep track of which digits are significant.
Why is it important to keep track of significant digits?
It is important to keep track of significant digits because your results should not imply more certainty than your measure quantities justify. This is especially important when you use a calculator. Calculators usually report results with far more significant digits – than you data warrant. There are three rules for reporting significant digits in calculated answers.
Complete the rules for reporting significant digits in calculations in the table below:
Table 1 Rules for Reporting Significant Digits in Calculations
Rule 1 Multiplying and Dividing
The value with the fewest number of significant digits, going into the calculation, determines the number of significant digits that you should report in your answetr.
Rule 2 Adding and Subtracting
The value with the fewest number of decimal places going into the calculation determines the number of decimal places that you should report in your answer.
Rule 3 Rounding
To get the appropriate number of significant digits (rule 1)or decimal places (rule 2), you may need to round your answer.
If your answer ends in a number that is greater than 5, increase the preceding digit by 1. For example, 2.346 can be rounded to 2.35.
If your answer ends in a number that is less than 5, eave the preceding number unchanged. For example. 5.73 can be left as 5.7.
If your answer ends with 5, increase the preceding number by 1 if it is odd. Leave the preceding number unchanged if it is even. For example, 18.35 can be rounded to 18.4, but 18.25 is rounded to 18.2.
Sample Problem Reporting Volume Using Significant Digits
A student measured a regularly shaped sample o f iron and found it to be 6.78 cm long, 3.906 cm wide and 11 cm tall. Determine its volume to the correct number of significant digits.
Length = 6.78 cm (3 sig digits)
Width = 3.906 cm (4 sig digits)
Height = 11 cm (2 sig digits)
Solution
V = l x w x h = 6.78 cm x 3.906 cm x 11 cm
= 291.30948 cm3
The value 11 cm has the smallest number of significant digits, 2. Thus your answer can only have two significant digits. In order to have only two significant digits, you need to put your answer into scientific notation: v = 2.2 x 102 cm
Sample Problem Reporting Mass Using Significant Digits
Suppose that you measure the masses of four objects as 12.5 g, 145.67 g, 79.0 g, and 38.438 g.
What is the total mass of the four objects ?
Strategy:
• Add your masses together, aligning them at the decimal place.
• Underline the estimated (farthest right) digit in each value. This will help you keep track of the number of estimated digits in your final answer.
• In the question, two values have the fewest decimal places: 12.5 and 79.0.
• Round your answer so that it has only one decimal place.
12 .5 145.67 79.0 38.438 275.608 ∴ total mass of the objects is 275.6 g
(this makes sense because the answer has one decimal place, the same as the value in the question with the fewest decimal places
Complete Practice Problems # 3 a – g, p. 22 text)
Date: Name: Class:
Reinforcement Reporting Significant Digits in Calculations
Chapter 1
BLM 1-2
Goal
Reinforce your understanding of significant digits.
Procedure
1. State the number of significant figures in each of the following:
(a) 3570
(b) 17.505
(c) 41.400
(d) 0.51
(e) 0.000 572
(f) 0.009 00
(g) 41.50 × 10-4
(h) 0.007 160 × 105
(i) 1.234 00 × 108
(j) 0.000 410 0 × 107
2. Perform the following operations and give the answer to the correct number of significant digits.
(a) 15.1 + 75.32
(b) 178.904 56 – 125.805
(c) 4.55 × 10-5 – 3.1 x 10-5
(d) 0.000 159 + 4.0074
(e) 1.805 × 104 + 5.89 × 104
(f) 0.000 817 - 0.000 048 1
(g) 8.166 × 105 – 7.819 × 105
(h) 45.128 + 8.501 87 – 42.18
(i) 5.677 × 10-6 + 7.785 × 10-6
(j) 8.75 × 10-9 + 6.1157 × 10-9
(k) 1.99 ÷ 3.1
(l) 1200.0 ÷ 3.0
(m) 5.32 x 10-4 ÷ 4.218 × 10-8
(n) 45.32 x 2.3
(o) 0.024 00 ÷ 6.000
(p) 12.4 x 0.30
(q) (5.50 x 108) × (4 x 105)
(r) 7.4 ÷ 3
(s) 4.75 × 5
(t) 2.5 × 6.700 ÷ 0.891
To
Solutions to BLM1-2 Significant Digits
1.
(a) 3 or 4
(b) 3
(c) 5
(d) 2
(e) 3
(f) 3
(g) 4
(h) 4
(i) of 6
(j) 4
2. (a) 90.4
(b) 53.0991
(c) 1.5 ﾗ10 −5
(d) 4.0076
(e) 7.70 ﾗ10 4
(f) 0.000 770
(g) 0.347 ﾗ10 5
(h) 11.45
(i) 1.346 ﾗ10 −5
(j) 1.49 ﾗ10 −8
(k) 0.64
(l) 4.0 ﾗ10 2
(m) 1.26 ﾗ10 4
(n) 1.0 ﾗ10 2
(o) 4.000 ﾗ10 −3
(p) 3.7
(q) 2 ﾗ10 14
(r) 2
(s) 2 ﾗ10 1
(t) 19
CHAPTER 1 BLM ANSWER KEY
2.1 THE ATOMIC THEORY OF MATTER (Text p.p. 34 – 41) In 1809, John Dalton described atoms as solid, indestructible particles that make up all matter. Dalton’s model of the atom was often referred to as the _ - _ model.
Dalton’s concept of the atom was one of several ideas in his atomic theory of matter. Dalton’s atomic theory of matter (1809) has four main points:
• ________________________________________________________________________
________________________________________________________________________
• ________________________________________________________________________
________________________________________________________________________
• ________________________________________________________________________
________________________________________________________________________
• ________________________________________________________________________
________________________________________________________________________
________________________________________________________________________
• Law of Conservation of Mass: during a chemical reaction, the total mass of the substances involved does not change (mass of __________________________ equals the mass of the ______________________________
• Law of Definite Proportions: Elements always combine to form compounds in fixed _____________________________by mass. For example, pure water always contains the elements hydrogen and oxygen combine in the following proportions: 11 % hydrogen, and 89% oxygen
Chemistry Background
What is a physical property ?
Physical properties—colour, density, boiling point, melting point—are properties of matter that are observable without changing one type of matter into another type of matter. What is a chemical property? Chemical properties—reactivity (with oxygen, with acid, or with other substances), toxicity—are properties that are observable when one type of matter changes into a different kind of matter of different properties.
THE MODERN VIEW OF THE ATOM An is the smallest particle of an element that still retains the identity and properties of the element. For example, the smallest particle of the writing material in a pencil is a carbon atom. Pencil “lead” is actually a substance called graphite. Graphite is a form of the element carbon. An average atom is about 10 – 10m in diameter. Atoms themselves are made up of even smaller particles. These subatomic particles are , , and . Protons and neutrons cluster together to form the central core, or , of an atom. Fast moving electrons occupy the space that surrounds the nucleus of the atom. As their name implies, subatomic particles are associated with electrical charges as shown in the table below:
Subatomic Particle
Charge
Symbol
Mass (g)
Radius (m)
Electron
Proton
Neutron 0
N0
1.67 x 10- 24
10 – 15
EXPRESSING THE MASS OF SUBATOMIC PARTICLES Subatomic particles are incredibly small. For example one proton or a neutron has a mass of g. It is inconvenient to measure mass of subatomic particles using units such as grams. Instead, chemists use a unit called an atomic mass unit (symbol ___). A proton has a mass of about 1 u, that is equal to 1.66 x 10 – 24 g. A proton is about 1837 times more massive than an electron. From the table above, the mass of an electron is 9.02 x 10 – 28. This value is so small that scientists consider the mass of an electron to be approximately equal to zero. Thus, electrons are not taken into account when calculating the mass of an atom.
EXPRESSING THE MASS IF SUBATOMIC SUBSTANCES
• Atomic Mass Unit
The Nucleus of an Atom
• Atomic Number (_______)
• Mass Number (______)
Information about an element’s protons and neutrons is often summarized using the chemical notation shown in Figure 2.3. Mass # A
X
Atomic # Z
In the figure above, which letter represents the atomic symbol of the element ? (X)
In the following example, identify the mass number and the atomic number
19 F 9
What do these numbers indicate ?
The mass umber (the superscript 19) indicates that fluorine has a total of 19 protons and neutrons. The atomic number (the subscript 9) indicates that fluorine has 9 protons.
How can you determine the number of protons in the nucleus of an atom?
To find the number of neutrons, subtract the atomic number from the mass number.
Number of Neutrons = mass # - atomic #
Or
Number of Neutrons = A - Z
How many neutrons would this isotope of fluorine have?
Number of Neutrons = A - Z = 19 - 9 = 10
PRACTICE PROBLEMS
1. Complete the missing values or chemical notations about the atoms from the table on p. 37 of your textbook below.
in a. b. c.
d. e f.
g h i
j k l
m n o
p. q. r.
s. t. u.
v. w. x.
When chemists refer symbolically to an element, they will often leave out the atomic number. For example, for oxygen, they would write 16 O or O – 16.
USING THE ATOMIC NUMBER TO INFER THE NUMBER OF ELECTRONS
The atomic number and mass number do not give you direct information about the number of neutrons in an element. They do not give you the number of electrons either. You can infer the number of electrons, however, from the atomic number. The atoms of each element are electrically neutral. This means that the positive charges (protons) and negative charges (electrons) must balance one another. Therefore, in a neutral atom of any element, the number of protons is equal to the number of electrons. For example, a neutral hydrogen atom contains one proton, so it must contain one electron. A neutral oxygen atom contains six protons, so it must contain six electrons.
ISOTOPES AND ATOMIC MASS All neutral atoms of the same element contain the same number of protons and, therefore, the same number of electrons. The number of neutrons can vary, however. For example, most of the oxygen atoms in nature have eight neutrons in their atomic nuclei. In other words, most oxygen atoms have a mass number of 16 (8 protons + 8 neutrons). As you can see in the figure below there are also two other naturally occurring forms of oxygen. One of these has nine neutrons, so A = 17, and the other has ten neutrons, so A = 18. These different forms of oxygen are called . Isotopes are ______
Fig. 2.4
p. 38 Text
The isotopes of an element have very chemical properties because they have the same number of protons and electrons. They differ in , however, because they have a different number of neutrons. Some isotopes are more unstable than others. Their nuclei are more likely to decay, releasing energy and subatomic particles. This process, called ______, happens spontaneously. All uranium isotopes, for example, have unstable nuclei. They are called radioisotopes. May isotopes are not radioisotopes. Oxygen’s three naturally occurring isotopes, for example, are stable. In contrast, chemists have successfully synthesized ten other isotopes of oxygen, all of which are unstable radioisotopes. Radioisotopes decay because their nuclei are unstable. The time it takes for a nucleus to decay varies greatly. For example, it takes billions of years for only half of the nucleus of naturally occurring uranium – 238 to decay. The nuclei of other radioisotopes, mainly those synthesized by scientists, decay much more rapidly. The nuclei of some isotopes, such as sodium – 22, take about 20 years to decay. For calcium – 47, this decay occurs in a mater of days. The nuclei of most radioisotopes decay so quickly, however, that the radioisotopes exist for mere fractions of a second. Atomic Weights
This is also called the “Average Atomic Mass”. It is the weighted average of all naturally occurring isotopes of an element. A.W. = (% 1) (m 1) + (% 2) (m 2) + (% 3) (m 3) . . . . . . . . 100
where % = the percent occurrence of the isotope in nature where m = the mass of the isotope
Example 1 Sulfur has three isotopes which have the following relative abundances: S – 32 = 95% S – 33 = 0.8 % S – 34 = 4.2 % Calculate the atomic weight of sulfur.
ELECTRONS IN ATOMS
What is the importance of the electron to the atom ? Recall that electrons occupy the space surrounding the nucleus therefore they are the first subatomic particles that are likely to interact when atoms come near one another. It is the and ________ of electrons that determine how an atom will react, if at all, with other atoms. ______are responsible for the chemical properties of elements.
REVISITNG THE ATOMIC THEORY John Dalton did not even know about subatomic particles when he developed his atomic theory. Even so, the modern atomic theory, shown below, retains many of Dalton’s ideas, with only a few modifications. Not all chemists believed that Dalton’s atoms existed. Other scientists considered atoms to be a valuable idea for understanding matter and its behavior. They did not , however, believe that atoms had any physical reality. The discovery of electrons (and later, the other subatomic particles) finally convinced scientists that atoms were more than simply an idea. Atoms, they realized, must be ______. The atomic theory is a landmark achievement in the history of chemistry. It shaped the way that all scientists, especially chemists, think about matter.
Table The modern Atomic Theory
p. 39 Text
SECTION REVIEW
1. For the table below, fill in the missing information in the appropriate space provided.
ELEMENT ATOMIC NUNMBER
MASS NUMBER
NUMBER OF PROTONS
NUMBER
OF ELECTRONS
NUMBER OF NEUTRONS
(a)
(b)
108
(c)
47
(d)
(e)
(f)
(g)
33
(h)
42
(i)
35
(j)
(k)
(l)
45
(m)
79
179
(n)
(o)
(p)
(q)
(r)
(s)
(t)
50
69
1. a. b. c. d.
e f. g. h.
i. j. k. l.
m. n. o. p.
q. r. s. t.
2. Explain the difference between an isotope and a radioisotope. Provide an example other than oxygen to support your answer.
3. Examine the information presented in the following pairs
3 3 14 16 19 18 H and He C and N F and F 1 2 6 7 9 9
(a) For each pair, do both members have the same number of protons ? electrons ? neutrons?
(b) Which pair or pairs consist of atoms that have the same value for Z? Which consists of atoms that have the same value for A ?
4. Compare Dalton’s atomic theory with the modern atomic theory. Explain why scientists modified Dalton’s theory.
5. Fill in the following table:
Protons
Electrons
Neutrons
P – 31
Al – 27
H-1 (protium)
H – 2 (deuterium)
NOTES
NOTES
6.
# 15, & 17 FROM Sch3a workbook 2000
Unit 2
7.
EXPANDING THE MODEL OF THE ATOM (p.p. 656 – 659)
Emission Spectra
Each wavelength of visible light is associated with a ______. When white light is shone through a gas discharge tube, it produces a bright line ______. A bright line spectrum is a series of narrow lines having specific colour ( _____), separated by colourless space. Each element produces a different and characteristic colour spectrum. You can see several examples of these spectra in the figure below. Scientists who were studying the atom, needed to take emission spectra into account.
Fig. D1
p. 656 text
Bohr’s Model of the Atom In 1913, Danish physicist, Niels Bohr developed a model of the atom that explained hydrogen emission spectrum. In Bohr’s model electrons ______ the nucleus in the same way that the Earth orbits the sun, as shown in the figure below.
Fig. D-2
p. 656 Text
The following three points of Bohr’s theory help explain hydrogen’s emission spectrum:
• In Bohr’s model, atoms have specific allowable ______levels. He called these energy levels stationary states. Each of these levels corresponds to a fixed, circular orbit around the nucleus.
• An atom does not give off energy while its electrons are in a stationary state
• An atom changes states by giving off, or absorbing a quantity of light energy exactly equal to the ______in energy between the two stationary states.
Bohr’s model was revolutionary because he proposed that the energy absorbed or emitted by an atom needed to have ______ ______. The energy change was ______, rather than continuous. When something is quantized, it is limited to discrete amounts or multiples of discrete amounts. Two great scientists paved the way for this surprising idea. German physicist Max Plank had already proposed that light was quantized, meaning that it exists in packets. Building on this idea, Einstein proposed that light could behave as particles, which he called photons. The energy associated with the light in a line spectrum corresponds with the change in energy of an electron as it moves _______or _______________ an energy level. For example, when electrons in hydrogen atoms that have been excited to the third energy level and subsequently drop to the second energy level, they emit light that has a specific wavelength. They emit photons of _____________ light that have a wavelength of 656.5 nm. These photons cause the red line spectrum for hydrogen seen in Figure D1 above. Why A New Model ? Bohr’s model worked fine at predicitng the line spectrum for hydrogen or for ions with only one electron e.g. He +, Li 2+ , Be 3+ etc. The model failed, however, when it was applied to the emission spectra of atoms that had more than one electron. Bohr’s model needed to be modified because it was too simple to explain the experimental evidence. Sublevels The spectra of many electron atoms suggested that a more complex structure was needed. Notice that in Figure D1, that spectra for these more complex atoms have groups of lines close together. The groups of lines are separated by spaces. The large spaces represent energy differences between ______ ______, while the smaller spaces represent energy ______within levels. If the electrons are changing energy within the levels, this suggests that there are sublevels within each level., each with its own slightly different energy. The idea of different energy levels 1, 2, 3, 4, . . . remains, but each energy level is split up into sublevels called _____, _____, _____, _____. Examine the table below to see how electrons are arranged in sublevels for energy levels 1 to 4.
Table D1
p. 657 text Each sublevel has its own energy. Examine the figure on the following page to see the relative energies of the different levels and sublevels. But where would you find electrons in this new model ? How are the sublevels oriented in space ?
Particles with Wave Like Properties
In 1924, Louis de Broglie suggested that all matter had -like properties. De Broglie developed an equation that allowed him to calculate the wavelength associated with any object. Objects we can see and interact with have calculated wavelengths that are smaller than electrons. Their wavelengths are so tiny compared to their size, that they do not have any measurable effect on the motion of the objects
For very tiny moving particles such as electrons, however, the wavelength becomes very significant. De Broglie’s theory was proven by experiment when streams of electrons produced diffraction patterns similar to electromagnetic radiation that was already known to travel in waves.
Orbitals
In 1920, Erwin Schrodinger used de Broglie’s idea that matter has wavelike properties. Schrodinger proposed the quantum mechanical model of the atom. In this new model, he abandoned the notion of the electron as a small particle orbiting the nucleus. Instead, he described the behaviour of electrons in terms of wave functions.
This imprecise nature of Schrodinger’s model was supported shortly afterwards by principle proposed by Heisenburg. Heisneburg demonstrated that it is impossible to know both an electron’s path and its exact location. Therefore it is impossible to know the exact location of an electron at any time, without some degree of uncertainty. Heisenburg’s uncertainty principle shows that you can never know both the ______and the ______ of an electron beyond a certain measure of precision. Heisneburg also showed that if you could know either the velocity or position precisely, then the other property would be uncertain. Therefore, when talking about where electrons are found, you can only talk in terms of
______ ______.
Schrodinger used a mathematical wave equation to define the probability of finding an electron within an atom. There are multiple solutions to this wave equation, and Schrodinger called these solutions ______. Each solution provides information about the energy and location of an electron within an atom. Each orbital has a specific energy associated with it, and each contains information about where inside the atom, the electrons would spend most of their time.
These regions of space where an electron is most likely to be found are called ______. In solving the wave equation, three unknowns were encountered (a fourth was added later). These unknowns are called the ______ ______and describe different properties of the electron.
The Four Quantum Numbers i. The Principle Quantum Number (_____)
• Indicates the ______an electron has
• Tells us which energy ______the electron is in
• Can have values of n = ______
• The higher the ‘n’ value, the ______the energy
• Correspond to Bohr’s orbits
2. The ______Quantum Number (_____)
• Indicates the ______of the orbital
• It can have values of L = ______ ______
• n- value possible ‘L’ values
1
2
3
4.
Note: the total number of possible ‘L’ values is the same as the ‘n’ value L- VALUE
0
1
2
3
4
SHAPE
SPEHRICAL
DUMBBELL
DOUBLE DUMBBELL
??
??
NAME
____– orbital
____– orbital
____– orbital
____– orbital
____– orbital
3. The _____________________Quantum Number (_____)
• indicates the _____________________of the orbitals about the nucleus.
• The number of values indicates the number of orbitals that exist with a given
• M L can have the values of ____________________________________ L- VALUE
0
1
2
3
4
SHAPE
SPEHRICAL
DUMBBELL
DOUBLE DUMBBELL
??
??
NAME
S – orbital
P – orbital
D – orbital
F – oribital
G - orbital
3. The Magnetic Quantum Number (M L)
• indicates the orientation of the orbitals about the nucleus.
• The number of values indicates the number of orbitals that exist with a given
• M L can have the values of - L . . . 0 . . . + L
L
0
1
2
3
4
5
NAME
S
P
D
F
G
H
M L
0
-1, 0, 1
-2, -1, 0, 1, 2
-3, -2, -1, 0, 1, 2, 3
NUMBER OF POSSIBLE ORBITALS
1
3
5
7
4. _________________Quantum Number (_____)
• Related to the _________________ of an electron
• All electrons spin on their axis. The two different spins are assigned the value of
________________________
• indicates that each orbital can hold a maximum of ________ electrons. If there are two electrons in an orbital, they must have opposite spins.
n
L
NAME
M L # of
Orbitals
M S # of
Electrons Total # of
Electrons
1
2
2
3
3
3
4
4
4
4
The ‘n’ Rules
• The nimber of different shapes (or ) of orbitals in a given energy level =
• The total number of orbitals in a given energy level =
• The total number of electrons in a given energy level =
Shapes of Orbitals The shapes of the probability graphs of Schrodinger’s wave functions are the shapes of the orbitals in which electrons reside in atoms. You can visualize orbitals as electron clouds. The shape of each cloud is based on probability – it tells where the electron spends most of its time. Examine the figure below. The ‘s’ orbitals are spherical in shape. Each of these spherical shells contains two electrons. There are three ‘p’ orbitals for each sublevel, each with a capacity for two electrons. Each ‘p’ orbital is shaped something like a dumbbell. For each subleve, the ‘p’ orbitals are oriented along the x, y, and z axes. The d’ and ‘f’ orbitals are quite complex in shape. Each ‘d’ sublevel contains 5 ‘d’ orbitals and each ‘f’ sublevel contains 7 ‘f’ orbitals.
S & p orbitals
SCH3a 2000 workbook
Quantum unit
D orbital diagram
Sch3a workbook 2000
f orbital diagram
Sch3a workbook 2000
Filling Orbitals
Each orbital can hold a maximum of ________ electrons. A hypothesis to suggest why this is true suggests that electrons spin around their own axes as they move around the nucleus, generating magnetic fields. They can spin in a positive direction or a negative direction . Electrons with opposite spins attract each other. This attraction partially counteracts the repulsion between two negatively charge electrons. In 1925, Pauli proposed that only two electrons of opposite spin could occupy an oribital.
The Pauli Exclusion Formula - states that ________________________________________________________________________
How do electrons fill orbitals within atoms? They do so to minimize the potential energy of the atom.
1. The maximum number of electrons that any orbital can hold is _____________
2. They will fill the ________________________energy level first. The 1s orbital will fill before the 2s orbital, which fills before the 2 p orbitals and so on.
1. When occupying two or more orbitals with the same energy (for example any of the 3 ‘p’ orbitals), electrons will half fill each orbital until all are half filled before adding a second electron to each one. This is called ________________________Rule. You can show how electrons fill orbitals using superscripts in a notation called the electron configuration of the atom. For example, a boron atom contains 5 electrons. The electron configuration for Boron would be 1s 2 2s 2 2p 1. The electron configuration for Nitrogen is 1s 2 2s 2 2p 3.
First half of Periodic Table & Their Quantum numbers
SCH3A Workbook 2000
Electron Configurations
2nd half of Periodic Table & Their Quantum numbers
SCH3A Workbook 2000
1. 21 Sc
2. 23 V
3. 30 Zn
4. 35 Br
5. 50 Sn
6. 57 La
7. 56 Ce
8. 64 Gd
9. 50 Hg
10. 84 Po
A Schematic Energy Level diagram of a many electron atom
Whole page
Sch3a 2000
Quantum Unit
Whole page
Quantum Review 1
Quantum Review 1 page 2
Notes
Notes
ATOMS, ELEMENTS, AND THE PERIODIC TABLE (Text p.p. 40 – 47) By the mid 1800’s, there were 65 known elements. Chemists studied these elements intensively and recorded detailed information about their reactivity and the masses of their atoms. Some chemists began to recognize patterns and behaviour of many of these elements.
Fig 2.5
p. 40
Other sets of elements display similar trends in their properties and behaviour. For example, oxygen (O), sulfur (S), selenium (Se), and tellurium (Te) share similar properties. The same is true for fluorine (F), chlorine (Cl), bromine (Br), and iodine (I). These similarities prompted chemists to search for a fundamental property that could be used to organize all elements. One chemist, Dmitri Mendeleev (1834 – 19907) sequenced the known elements in order of increasing atomic mass. The result was a table of the elements, organized so that elements with similar properties were arranged in the same ________________________. Because Mendeleev’s arrangement highlighted the periodic (repeating) patterns of properties, it was called a periodic table. The term ________________________means “repeating in an identifiable pattern.
The modern periodic table is a modification of the arrangement first proposed by Mendeleev. Instead, of organizing the elements in order of increasing atomic _____________, the modern periodic table organizes elements according to atomic ________________________. According to the Periodic Law, the chemical and physical properties of the elements repeat in a regular, periodic pattern when they are arranged according to their atomic number.
p. 41 text PRACTICE PROBLEMS
2. Identify the name and symbol of the elements in the following location of the periodic table:
a. Group 14 (IVA), Period 2
b. Group 11 (IB), Period 4
c. Group 18 (VIIIA), Period 6
d. Group 1 (IA), Period 3
e. Group 12 (IIB), Period 5
f. Group 2 (IIA), Period 4
g. Group 17, (VIIA), Period 5
h. Group 13 (IIA), Period 3
_____________________________
_____________________________
_____________________________
_____________________________
_____________________________
_____________________________
_____________________________
_____________________________
ELECTRONS AND THE PERIODIC TABLE
We have seen how the periodic table organizes elements so that those with similar properties are in the same group. You have also seen how the periodic table shows a clear distinction among metals, non-metals, and metalloids. Other details of the organization of the periodic table may seem strange. Why, for example, are there different numbers of elements in the periods.
The reason for this, and other details of the periodic table’s organization, involves the number and arrangement of ________________________in atoms of each element.
ELECTRONS AND ENERGY LEVELS Electrons cannot move haphazardly. Their movement around an atomic nucleus is restricted to fixed regions of space. These regions of space are three dimensional. The following figure shows a representation of these regions.
Fig 2.8
p. 43 text
These regions are not solid, but rather are volumes of space in which electrons may be found. They are referred to as energy shells or energy levels. An electron that is moving in a lower energy level is found ________________________to the nucleus. It would have less energy than it would if it were moving in a higher energy level. There is a limit to the number of electrons that can occupy each energy level. For example, a maximum of _______________ electrons can occupy the first energy level. A maximum of ________________________electrons can occupy the second energy level. The periodic trends that result from organizing the elements by their atomic number are linked to the way in which electrons occupy and fill energy levels.
Text
Fig 2.9
p. 44
As shown in the figure above, a common way to show the arrangement of electron in atoms is to draw circles around the atomic symbol. Each circle represents an energy level. Dots represent electrons that occupy each energy level. This kind of diagram is called a Bohr –Rutherford diagram. Figure 2.9 B shows that the first energy level is full when _____________ electrons occupy it. Only two elements have two or fewer electrons: hydrogen and helium. Hydrogen has one electron, and helium has two. These elements, with their electrons in the first energy level, make up Period _______ of the periodic table. As you can see in Fig. 2.9 C, Period 2 elements have two energy levels. The second energy level is full when ________electrons occupy it. Neon, with a total of ten electrons, has its first and second energy levels filled. Notice how the second energy level fills with electrons as you move across the period from lithium to fluorine.
PATTERNS BASED ON ENERGY LEVELS AND ELECTRON ARRANGEMENTS The structure of the periodic table is closely related to ________________________ levels and the arrangement of electrons. Two important patterns result from this relationship. One involves periods, and the other involves groups. The Period Related Pattern
As you can see in Figure 2.9, elements in Period 1 have electrons in one energy level. Elements in Period 2 have electrons in two energy levels. This pattern applies to all seven periods. An element’s period number is the same as the number of ________________________levels that the electrons in the atom occupy. Thus, you could predict that Period 5 elements occupy 5 different energy levels. What about the inner transition elements – the elements that are below the periodic table? Figure 2.10 shows how this pattern applies to them. Elements 58 through 71 belong in Period 6, so their electrons occupy six different energy levels. Elements 90 through 103 belong in Period 7, so their electrons occupy seven different energy levels. Chemists and chemical technologists tend to use only a few of the inner transition elements on a regular basis. Thus it is more convenient to place all inner transition elements below the periodic table.
The Group Related Pattern The second pattern emerges when you consider the electron arrangements in the main- group elements: the elements in Group 1 (1A), 2 (2A) and 13 (3A) to 18 (8A). All the elements in each main group have the same number of electrons in their highest (outer) energy level. The electrons that occupy the outermost energy level are called ________________________ electrons. It is these valence electrons that are involved when atoms form compounds. In other words, ________________________ ________________________are responsible for the chemical behaviour of elements. You can infer the number of valence electrons in any main group element from its group number. For example, Group 1 (1A) elements have _____ valence electron. Group 2 (2A) elements have _____ valence electrons. For elements in Group 13 (3A) to 18 (18A), the number of valence electrons is the same as the second digit in the current numbering system. It is the same as the number only in the old numbering system (in brackets). For example, Group 15 (5A) have five valence electrons while Group 17 (7A) have seven valence electrons.
USING LEWIS STRUCTURES TO REPRESENT VALENCE ELECTRONS It is time consuming to draw electron arrangements using Bohr-Rughterford diagrams. It is much simpler to draw Lewis structures to represent elements and their valence electrons. To draw a Lewis structure, you replace the nucleus and the inner energy levels of an atom with its atomic symbol. Then you place dots around the atomic symbol to represent the valence electrons. The order you place the first four dots is up to you, however, place a dot on each side of the symbol before doubling up on any dot.
Fig 2.11
Text
Practice Problems
2. On the blank periodic table, sketch the electron arrangements for the first 20 elements using Lewis structures.
3. Use the periodic table to draw Lewis structures for the following elements: barium (Ba), gallium (Ga), tin (Sn), bismuth (B), iodine (I), cesium (Cs), krypton (Kr) and xenon (Xe)
4. Identify the number of valence electrons in the outer energy levels of the following elements:
a. chlorine f. lead
b. helium g. antimony
c. indium h. selenium
d. strontium i. arsenic
e. rubidium j. xenon NOTES
THE SIGNIFICANCE OF A FULL OUTER ENERGY LEVEL The ________________________ ________________________in Group 18 (VIIIA) are the only elements that exist as individual atoms in nature. They are extremely ________________________. They do not naturally form compounds with other atoms. What is it about noble gases that explains their behaviour ? Recall that chemical reactivity is determined by ________________________electrons. Thus, there must be something about the arrangement of the electrons in noble gases that explains how unreactive they are. All the noble gases have outer energy levels that are completely ________________________with the maximum number of electrons. Helium has a full outer energy level of ______ valence electrons. The other noble gasses have ______ valence electrons in the outer energy level. Chemists reason that having a full outer energy level must be a very ________________________electron arrangement. What does this stability mean ? It means that a full outer energy level is unlikely to change. Scientists have observed that, in nature, situations or systems of lower energy are favored over situations of higher energy. For example, a book on a high shelf has more potential energy than a book on a lower shelf. If you move the book to the ground, it has low potential energy, that is, it is more stable. When atoms have eight electrons in their outermost energy level, (or two for hydrogen and helium), chemists say they have a stable ________________________. An octet is a very stable electron arrangement. Summary
You have seen that the structure of the periodic table is directly related to energy levels and arrangement of electrons. The patterns that emerge from this relationship enable you to predict the number of valence electrons for any main group element. They also enable you to predict the number of energy levels that an element’s electrons occupy.
Orbital Box Notation Using this method, each orbital is shown as a box with ________________________ representing the electrons. If there are two electrons in the same orbital box, then they must have opposite spins. This is shown by having the arrows facing in different _____________________.
Element Electron Configuration Orbital Box
4 Be
6 C
7 N
Hund’s Rule _______________________________________________________________________________________________________________________
8 O
9 F
21 Sc
50 Sn
Ground State Electron Configurations with this arrangement, the electrons are in the ________________________possible orbitals. This is the order given by the Periodic Table.
Excited State Configurations In order to obtain these configurations, energy must be added to the atom. When energy is added, one or more electrons is moved to a higher energy orbital. To which orbital the electron gets moved depends on the amount of energy that is added.
12 Mg Ground State Electron Configuration Excited State (one of many)
Electron Configurations and Valences Atoms react with other atoms to become more ________________________ (to have a more stable electron arrangement). It is the arrangement of the atom’s outermost electron that determine its chemical properties. To become more stable, an atom can either _______________ or ___________________ electrons. If an atom is giving away electrons, then the ones with the highest ______ value always go first. Note that the rule for taking electrons away does not follow the rule for adding electrons to an atom. Stable Arrangements There are several different arrangements which can make an atom more stable:
1. The most stable arrangement is achieved by having a completely filled energy level.
Examples:
2. The second most stable configuration involves having fully occupied o completely empty energy levels. A sub-energy level refers to all the orbitals of one-kind at that energy level.
Examples:
3. The third most stable configuration is achieved by having half-filled sub-energy levels Examples:
Valences Predict the most likely valences for the following:
1. 7 N
2. 16 S
3. 30 Zn
4. 50 Sn
Explain why:
i. Fe (# 26) has a (+2) and a (+3) valence. Which is more stable ?
ii. Rh (# 45) has a (+2) and a (+4) valence.
iii. Cerium (# 58) has a (+3 and a (+4) valence
Exceptions to the Predicted Electronic Configurations There are two families whose electron configurations are different than what would be predicted. They are the families VIB and IB. The elements in these families spontaneously rearrange their electrons to make themselves more stable. To do this, they promote an electron from an ‘s’ orbital and excite it to a ‘d’ orbital.
Examples
24 Cr
42 Mo
29 Cu
47 Ag
Chemical Rections and the desire to become more stable
Sch 3a Workbook 2000
Section Review
1. State the Periodic Law and provide at least two examples to illustrate its meaning.
2. Identify the Group number for each of these sets of elements. Then write the symbols for the elements within it:
• alkali metals
• noble gases
• halogens
• alkaline earth metals
3. Identify the elements that is described by the following information. Refer to a periodic table as necessary.
• It is a Group 14 (IIIA) metalloid in the third period
• It is a Group 15 (VA) metalloid in the fifth period
• It is the other metalloid in Group 15
• It is a halogen that exists in the liquid state at room temperature
4. What is the relationship between electron arrangement and the organization of the elements of the periodic table ?
5 How many valence electrons are there in an atom of each of these elements ?
(a) neon sodium magnesium bromine chlorine silicon sulfur helium strontium tin
(b) Present your answers from Part A in the form of Lewis structures.
(c) Classify each element from Part A as a metal, non-metal, or a metalloid
6. Compare and contrast noble gases with other elements.
7.
(a) Draw Lewis Structures for each of the following elements: lithium potassium magnesium aluminum carbon
(b) Which of these elements have the same number of energy levels ?
(c) Which have the same number of valence electrons
8. Identify the elements with the following electron distributions by orbits:
a. 2, 8, 8 b. 2, 8, 6 c. 2, 4 d. 2, 8, 8, 2
5. What is meant by a stable octet ?
6. How many atoms must each of the following elements gain to achieve a stable octet ?
a. fluorine b. oxygen c. nitrogen d. carbon
7. How many electrons must the following atoms lose in order to achieve a stable octet
a. silicon b. phosphorus c. sodium d. argon
8. How does the sign of the valence relate to the gain or loss of electrons ?
9. Explain why atoms form ions.
Notes
Notes
Notes
Notes
Notes
2.3
PERIODIC TRENDS INVOLVING THE SIZES AND ENERGY LEVELS OF ATOMS In Section 2.1, we learned that the size of a typical atom is about 10 – 10 m, however, we also know that the atoms of each element are distinctly different. For example, the atoms of different elements have different numbers of protons. This means of course, that they also have different numbers of electrons. You may predict that the size of an atom is related to the number of protons and electrons it has. Is there any evidence to support this prediction, If so, is there a pattern that could help you to predict the relative size of an atom for any element on the periodic table ? Chemists define, and measure, an atom’s size in terms of its radius. The radius of an atom is the distance from its nucleus to the approximate outer boundary of the cloud-like region of its electrons. This boundary is approximate because atoms are not solid spheres. They do not have an outer boundary.
Fig. 2.12
p. 49 Text
Inv 2-A
p.p. 50 & 51
Text
Inv 2-A
p.p. 50 & 51
Text
Trends for Atomic Size (Radius) There are two general trends for atomic size:
• As you go down each group in the periodic table, the size of an atom ___________________. This makes sense if you consider energy levels. As you go down a group, the valence electrons occupy an energy level that is further from the nucleus. Thus, the valence electrons experience less attraction for the nucleus. In addition, electrons in the inner energy levels block, or shield the valence electrons from the attraction of the nucleus. As a result, the total volume of the atom, and thus size, increase with each additional energy level.
• As you go across a period, the size of an atom ________________________. This trend might seen surprising at first, since the number of electrons increases as you go across a period. The size of the atom decreases, however, because the positive charge of the nucleus also increases across a period. As a result, the positive force exerted by the nucleus increases, pulling the outer electrons ________________________, thus reducing the atom’s total size
Fig 2.13
p. 552 Text
General Trends for Atomic Radius Graph
SCH 3A workbook 2000
Trends on he P.T.
Sch3a Workbook
Practice Problems
7. Using only their location in the periodic table, rank the atoms in each set by decreasing atomic size. Explain your rankings. (a) Mg Be Ba
(b) Ca Se Ga
(c) Br Rb Kr
(d) Se Br Ca
(e) Ba Sr Cs
(f) Se Br Cl
(g) Mg Ca Li
(h) Sr Te Se
(i) In Br I
(j) S Se O Notes
Trends for Ionization Energy A neutral atom contains an equal number of positive charges (protons), and negative charges (electrons). It takes ________________________to overcome the attractive force of the nucleus and pull an electron away from a neutral atom. The energy required to remove an electron from an atom is called ________________________energy. The particle that results when a neutral atom gains electrons or gives up electrons is called an ion. Thus, an ion is a ________________________particle. An ion that ________________________electrons becomes a negatively charged anion, while an ion that ________________________up electrons becomes a positively charged cation. The figure below shows the formation of ions for several elements. As you examine the diagrams, pay close attention to
• the energy level from which the electrons are gained or given up
• the charge of the ion that is formed when an atom gains or gives up electrons
• the arrangement of the electrons that remain after the electrons are gained or given up.
Fig . 2.14
p. 53 text
If you try to visualize the periodic table as a cylinder, (see Fig. 2.15 below) rather than a flat plane, you can begin to see the relationship between ion formation and the electron arrangement of the noble gases.
Fig 2.15
p. 54 Text
The metals that are the main-group elements tend to ______________ up electrons and form ions that have the same number of electrons as the nearest ________________________gas. Non-metals tend to _____________________ electrons and form ions that have the same number of electrons as the nearest noble gas. For example, when a sodium atom gives up its single valence electron, it becomes a ________________________charged sodium ion. Its outer electron arrangement is like ________________________outer electron arrangement. When a fluorine atom gains an electron, it becomes a ________________________charged ion with an outer electron arrangement like that of ________________________. Figure 2.15 can help you determine the charge on an ion. Count the number of groups an ion is from the nearest noble gas. That number is the ________________________on the ion. For example, aluminum is three groups away from neon, and so aluminum has a charge of + 3. Sulfur is two groups away from argon, thus sulfur has a charge of 2 -. Remember metals form positive ions (________________) while non-metals can form both positive ions (________________) or negative ions (________________). Negative ions are formed by non-metals when they combine with metals. One of the reasons that the periodic table is useful is that it can help us remember the types of ions formed by many of the representative elements. For example, except for hydrogen, the neutral atoms of Group 1A always lose one electron each when they react, thereby becoming ions with a charge of 1 +. Similarly, elements in Group II always lose two electrons to become ions with a charge of 2+. In Group III, the only important ion that we need to consider at this time is aluminum, which loses three electrons to form an ion with a charge of 3 +. Notice that the charges on the cations is the same as their traditional Group numbers on the periodic table. Although this generalization does not work for all metallic elements (it doesn’t work for the transition elements), it does help us remember what happens to the members of Group I, Group II, and aluminum when they react, Among the non-metals on the right side of the periodic table, we also find some useful generalizations. For example, when they combine with metals, the_____________________ (Group VIIA or Group 17) form ions with a 1 – charge, the non-metals in Group VIA or Group 16, form ions with a 2 – charge. Notice that the number of negative charges on the anion is equal to the number of spaces to the right that we have to move on the periodic table to get to a noble gas.
Littlediagram from new sjuce
Middle of page
The figure below shows the ionization energy that is required to remove one electron from the outer energy level of atoms of the main group elements This energy is called the ____________-ionization energy. It is measured in units of kJ / mol. As you can see, atoms that give up electrons easily have _______________ ionization energies. You would probably predict that the alkali metals of Group 1 (1A) would have low ionization energies. These elements are, in fact, extremely reactive because it takes so little energy to remove their single valence electron.
Fig. 2.16
p. 54 text
All elements, except _________________________, have more than one electron that can be removed. Therefore, they have more than one ionization energy. The energy that is need to remove a second electron is called the ________________________ionization energy. The energy needed to remove a third electron is called the ________________________ionization energy, and so on. Summarizing Trends for Ionization Energy
Upon examination of the periodic trend for ionization energy it is the opposite of the trend for atomic radius.
Fig. 2.17
Text p. 55
Bar Graph from SCH3a workbook 2000
Atomic Radius Vs I.E.
Although there are a few exceptions, there are two general trends for ionization energy:
• Ionization energy tends to ________________________down a group as electrons in the outer energy level are further from the positive force of the nucleus, thus they are easier to remove than electrons in lower energy levels.
• Ionization energy tends to ________________________across a period. As you go across a period, the attraction between the nucleus and the electrons in the outer energy level increases. Thus, ______________________ energy is required to pull an electron away from its atom. For this trend to be true, you would expect the noble gases to have the highest ionization energies, and as is shown in Figure 2.16, they do. Practice Problems
8. Using only a periodic table, rank the elements in each set by increasing ionization enegy. Explain your answers. Explanation
(a) Xe He Ar
(b) Sn In Sb
(c) Sr Ca Ba
(d) Kr Br K
(e) K Ca Rb
(f) Kr Br Rb
9. Using only a periodic table, identify the atom in each of the following pairs with the lower first ionization energy. (a) B O
(b) B In
(c) I F
(d) F N
(e) Ca K
(f) B Tl Trends for Electron Affinity Electron affinity is the change in ________________________that occurs when an electron is added to the outer energy level of an atom to form a negative ion. Figure 2.18 below gives the electron affinities of the main group of elements. If the ion that is formed by gaining an electron is stable, the electron affinity is expressed as a negative integer. The more ________________________ the ion, the higher is the negative integer for the electron affinity. Notice that fluorine has the highest electron affinity. This indicates that fluorine is very likely to be involved in chemical reactions. In fact, fluorine is the most ________________________of all elements. Metals have ___________ electron affinities. That is especially true for the Group 1 (1A) and 2 (2A) elements. Atoms of these elements form ________________________positive ions. A negative ion that is formed by elements of these groups is unstable. It breaks apart into a neutral atom and a free electron
Fig. 2.18
Text p. 57
Fig 2.19
p. 58 text
Periodic Trend for Electronegativity (p.p. 70 – 71) When two atoms form a bond, each atom attracts the other atom’s electrons in addition to its own. The electronegativity of an atom is a measure of the atom’s ability to ________________________ electrons in a chemical bond. _________ is used to symbolize electronegativity. There is a specific electronegativity associated with each element. While atomic radius, ionization energy, and electron affinity are all properties of single atoms, electronegativity is a property of atoms that are involved in chemical bonding.
Fig. 3.6
Text p. 71
The trend for electronegativity is the ________________________of the trend for atomic radius. In general, atomic size decreases from left to right across a period, electronegativity ________________________. This can be explained by looking at two factors. The number of ________________________ (which attract electrons) is increasing and at the same time, the number of energy levels (which shield the electrons from the attraction of the protons in the nucleus) remains the same. Thus electrons are pulled more tightly to the nucleus, resulting in a ________________________atomic size. In the second period, for example, lithium has the largest atomic radius and the lowest elecronegativity. As atomic size decreases across the second period, the electronegativity increases. Fluorine has the smallest atomic size in the third period, yet it has the largest electronegativity. Because noble gases do not usually participate in chemical bonding, their electronegativities are not given. Similarly, as atomic size increases down a group, electronegativity ________________________. As you move down a group, valence electrons are less strongly attracted to the nucleus because the number of energy levels between the nucleus and the valence electrons increases. In a compound, increasing energy levels between valence electrons and the nucleus mean that the nucleus attracts bonding pairs less strongly. For example, in Group 2 (IIA), berylliun has the smallest atomic radius and the largest electronegativity. As atomic size increases down a group, electronegativity ________________________. Section Review
1. How does your knowledge of electron arrangement and forces in atoms help you explain the following periodic trends ?
a. atomic radius b. ionization energy c. electron affinity
2. Using only their position on the periodic table, rank each of the following sets of elements in order of increasing atomic size. Explain your answer in each case.
a. Mg S Cl b. Al B In
c. Ne Ar Xe d. Rb Xe Te
e. P Na F f. O S N
3. Using only their position on the periodic table, rank each of the following sets of elements in order of decreasing ionization energy. Explain your answer in each case.
a. Cl Br I b. Ga Ge Se
c. K Ca Kr d. Na Li Cs
e. S Cl Br f. Cl Ar K
4. Which element in each of the following pairs will have the lower electron affinity ? Explain your answer.
a. K or Cl b. O or Li
c. S or Se d. Cs or F
Notes
Notes
Chapter 2 Review
p. 61
Notes
Chapter 2 Review
P. 62
Notes
Chapter 2 Review
P. 63
3.1
Classifying Chemical Compounds (p.p. 66 – 74) Most elements do not exist in nature in their pure form as ________________________. Gold, silver, and platinum are three metals that can be found in the earth’s crust as elements. They are called “precious metals” because their occurrence is so rare. Most other metals, and most other elements, are found in nature only as ________________________. There are only about 90 naturally occurring elements. In comparison, there are thousands upon thousands of different compounds in nature, and more are constantly being discovered. Because there are so many compounds, chemists have developed a classification system to organize them according to their properties, such as melting point, boiling point, hardness, conductivity and solubility.
Express Lab A Metal and a Compound Text p. 67
Complete the Express Lab along with Analysis Questions 1 – 4
Notes
Properties of Ionic and Covalent Compounds Based on their physical properties, compounds can be classified into two groups: ________________________compounds and ________________________compounds. Some of the properties of ionic and covalent compounds are summarized in the table below:
PROPERTY
IONIC COMPOUND
COVALENT COMPOOUNDS
State at room temperature
Melting Point
Electrical conductivity as a liquid
Yes
Solubility in water
Conducts electricity when dissolved in water
What is Bonding ? Chemical bonds are the forces that ________________________atoms to each other in compounds. Bonding involves the interaction between the ________________________ electrons of atoms. Usually the formation of a bond between two atoms creates a compound that is more ________________________than either of the two atoms on their own. The different properties of ionic and covalent compounds result from the manner in which chemical bonds form between atoms in these compounds. Atoms can either ________________________or ________________________electrons. When two atoms exchange electrons, one atom loses its valence electron(s) while the other atom gains the electron(s). This kind of bonding usually occurs between a metal and a non-metal. Recall, that metals have low ionization energies, while non-metals have high electron affinities. That is, metals tend to ________________________electrons, and non-metals tend to ________________________electrons. When atoms exchange electrons, they form an ________________________bond. Atoms can also share electrons. This kind of bond forms between two ________________________. It can also form between a metal and a non-metal when the metal has a high ionization energy. When atoms share electrons, they form a ________________________ bond. Each atom contributes one or more electrons to the covalent bond. The electrons are shared between the atoms.
Predicting Bond Type Using Electronegativity (p. 72 Text) You can use differences in electronegativities to decide whether the bond between atoms in ionic or covalent. When calculating electronegativity differences, the smaller value is always subtracted from the larger value so that the difference is always positive.
The Range of Electornegativity Differences When two atoms have electronegativity values that are identical, as in oxygen, they share the bonding their bonding pair of electrons equally between them in a pure covalent bond. When two atoms have electronegativity values that are different as in potassium fluoride, potassium gives up its valence electron to fluorine. Therefore, the bond is ________________________. It is not always clear whether atoms share electrons or transfer them. Atoms with different electronegativities can share electrons unequally without exchanging them. The figure below shows the range of electronegtivity differences. These values go from mostly covalent at 0.0 to mostly ionic at 3.3. Chemists consider bonds with an electornegativity difference that is greater than 1.7 to be ________________________and bonds with an electronegativity difference less than 1.7 to be ________________________.
Fig 3.8
p. 73 text
The table above shows you how to think of bonds having a percent ionic character, or a percent covalent character based on their electronegativity differences. When bonds have nearly 50 % ionic or covalent character, they have characteristics of both types of bonding. Covalent bonds between two atoms with different electronegativity values will have a certain degree of ionic character and are said to be ________________________covalent bonds.
Fig. 3.6 p. 71 text
Practice Problems
1. Determine the ΔEN for each bond shown. Indicate whether each bond I ionic or covalent
a. O - H
b. C - H
c. Mg - Cl
d. B - F
e. Cr - O
f. C - N
g. Na - I
h. Na - Br
Section Review
1. Name the typical properties of an ionic compound. Give two examples of ionic compounds.
2. Name the typical properties of a covalent compound. Give two examples of covalent compounds.
3. Describe and explain the periodic trend for electronegativity.
4. Based only on their position in the periodic table, arrange the elements in each set in increasing attraction for electrons in a bond.
5. Determine the ΔEN for each bond. Indicate whether the bond is ionic or covalent.
a. N - O b. Mn - O c. H - Cl d. Ca - Cl
Notes
Notes
3.2 IONIC AND COVALENT BONDINNG: THE OCTET RULE
The Octet Rule When atoms form bonds, they are often more stable. We know that noble gases are the most stable elements in the periodic table. They are extremely unreactive, and they tend not to form compounds. What the noble gases have in common is a _____________________ outer electron energy level. When an atom loses or gains electrons to achieve a filled outer electron energy level, the atom often becomes more stable. According to the octet rule, atoms bond in order to achieve an electron configuration that is the same as the electron configuration as a noble gas. When two atoms or ions have the same electron configuration, they are said to be isoelectronic. For example, Cl - is isoelctronic with Ar because both have 18 electrons and a filled outer energy level.
Ionic Bonding Sodium has a very low electornegativity, and chlorine has a high electronegativty. Therefore, when sodium and chlorine interact, sodium transfers its valence electron to chlorine. As shown in the figure below, sodium becomes Na + 1 and chlorine becomes Cl -.
Fig 3.9
p. 75 text
When neutral sodium loses its one valence electron to chlorine, the resulting Na + 1 cation has an en electron energy level that contains ___________ electrons. It is _____________________ with the noble gas neon. On the other hand, chlorine’s outer energy level has seven electrons. When chlorine gains sodium’s electron, it becomes an anion _____________________ with the noble gas argon.
The figure below shows how to represent the formation of an ionic bond using Lewis Structures. Thus, in an ionic bond, electrons are _____________________ from one atom to another so that they form oppositely charged ions. The strong force of attraction between oppositely charged ions is what holds them together.
Fig. 3.10
Text
Fig. 3.11 text Practice Problems
1. For each bond below, determine the ΔEN. Is the bond ionic or covalent ? a. Ca - O
b. K - Cl
c. K - F
d. Li - F
e. Li - Br
f. Ba - O
2. Draw Lewis structures to represent the formation of each bond in Question 1
Notes
COVALENT BONDING
We have just learned what happens when the electronegativity difference between two atoms is greater than 1.7. The atom with the lower electronegativity transfers its valence electron(s) to the atom with the higher electronegativity. The resulting ions have opposite charges. They are held together by a strong ionic bond. What happens if the electronegativity difference is zero. As an example, consider chlorine gas. Each chlorine atom has seven electrons in its outer energy level (valence shell). In order for chlorine to achieve the electron configuration of a noble gas according to the octet rule, it needs to gain one electron. When two chlorine atoms bond together, their electronegativity difference is zero and the electrons are equally attracted to each atom. Therefore, instead of transferring electrons, the two atoms each share one electron with each other. In other words, each atom contributes one electron to a covalent bond. A covalent bond consists of a pair of _____________________ electrons. Thus, each chlorine atom achieves a filled outer energy level, satisfying the _____________________ rule. The figure below shows how to represent a covalent bond with a Lewis structure.
Fig . 3.15 text
When two atoms of the same element form a bond, they share their electrons equally in a _____________________ covalent bond. Elements that bond to each other in this way are known as _____________________ molecules. When atoms such as carbon and hydrogen bond to each other, their electronegativities are so close that they share electrons almost equally. Carbon and hydrogen have an electronegativity difference of only 2.6 -= 2.2 = 0.4. In the figure below, you can see how one atom of hydrogen forms a covalent bond with four atoms of hydrogen. The compound methane, CH4 , is formed.
Fig 3.16 text
Each hydrogen shares one of its electrons with carbon. The carbon shares its four electrons with each hydrogen atom. Thus, each hydrogen achieves a filled outer energy level, as does carbon. (Recall that elements in the first period only need two electrons to fill their outer energy level). When analyzing Lewis structures that show covalent bonds, count the shared electrons as if they belong to each of the bonding atoms.
Practice Problems
1. Show the formation of a covalent bond between two atoms of each diatomic element. a. iodine
b. bromine
c. hydrogen
d. fluorine
2. Use Lewis structures to show the simplest way in which each pair of elements forms a covalent bond, according to the octet rule. a. hydrogen and oxygen
b. chlorine and oxygen
c. carbon and hydrogen
d. iodine and hydrogen
e. nitrogen and hydrogen
f. hydrogen and rubidium
Mutliple Covalent Bonds (p. 82 Text) In covalent bonding, atoms sometimes need to share two or three pairs of electrons, according to the octet rule. For example, consider the diatomic element oxygen. Each oxygen atom has six valence electrons in its outer shell, and therefore, each atom requires two additional electrons to achieve a stable octet. When two atoms form a bond by sharing two pairs of electrons, it results in a _____________________ covalent bond. When atoms form a bond by sharing three pairs of electrons, they form a _____________________ bond. An example of a triple bond is seen in diatomic nitrogen.
Figures 3.17, 3.18, 3.19 p. 82 Text
Note: the number of bonds each atom must form is equal to its ________________ Section Review
1. Use Lewis structures to show how each pair of elements forms an ionic bond.
a. magnesium and fluorine b. potassium and bromine
c. rubidium and chlorine d. calcium and oxygen
2. Use Lewis structures to show how each pair of elements forms covalent bonds.
a. one silicon atom and two oxygen atoms
b. one carbon atom, one hydrogen atom, and three chlorine atoms
c. two nitrogen atoms
d. two carbon atoms bonded together with three hydrogen atoms bonded to one carbon atom and one hydrogen atom and one oxygen atom bonded to the other carbon.
3. Use what you know about electronegativity differences to decide what kind of bond would form between each pair of elements.
4. In general, the farther away two elements are form each other on the periodic table, the more likely they are to participate in ionic bonding. Do you agree with this statement? Explain why or why not.
Notes
Polar Covalent Bonds: The “In-Between Bonds”
When two atoms have an elecronegativity difference hat is greater than 0.5 but less than 1.7, they are considered to be a polar covalent bond. In a polar covalent bond, the atoms have significantly different electronegativities. The electronegativity difference is not great enough, however, for the less electronegative atom to transfer its valence electrons to the other, more electronegative atom. The difference is great enough for the bonding electron pair to spend more time near the more electronegative atom. For example, the bond between oxygen and hydrogen in water has an electronegativity difference of 1.24. Because this value is less than 1.7 but greater than 0.5, the bond is a polar covalent bond. The oxygen attracts the electrons more strongly than the hydrogen. Therefore, oxygen has a slightly negative charge and the hydrogen has a slightly positive charge. These charges are not full charges but rather a partial charge symbolized by the symbols δ + and δ -.
Fig 3.24 & Fig 3.25
P. 86 Text
Practice Problems Predict whether each bond will be covalent, polar covalent, or ionic.
a. C – F
b. O – N
c. Cl – Cl
d. Cu – O
e. Si – H
f. Na – F
g. Fe – O
h. Mn – O
2. For each polar covalent bond in Question 1, indicate the location of the partial charges.
3. Arrange the bonds in each set in order of increasing polarity. (A completely polarized bond is an ionic bond)
a. H – Cl O – O Na – Cl
b. C – Cl Mg – Cl P – O N - N
Notes
Notes
Fig. 3.26
p. 87 text
Lab of some sort
Section Review
p. 94 text questions
Notes
Appendix 1 Periodicity and Bonding Questions
Review Perodicity & Bonding p. 1
Sch3a Worlbook
Appendix 1 Periodicity and Bonding Questions
Review Perodicity & Bonding p. 2
Sch3a Worlbook
3.4
WRITING CHEMICAL FORMULAS AND NAMING CHEMICAL COMPOUNDS
A chemical formula provides two important pieces of information:
• the elements that make up the compound
• the number of atoms of each element that are present in a compound
The order in which the elements are written also communicates important information. Convention, the less electronegative element is always listed first in the formula, the more electronegative element comes second. For example, the ionic compound that is formed from calcium and bromine is written CaBr 2. Calcium, a metal with a low electronegativity, is listed first. The subscript 2 after bromine indicates that there are 2 bromine atoms for every calcium atom.
Fig. 3.39 & 3.40
P. 95 Text
Then all of Unit 3 from Workbook, beginning with Nomenclature notes at the end of unit 3
CHAPTER 4: CLASSIFYING CHEMICAL REACTIONS: CHEMICALS IN BALANCE
4.1 Chemical Equations Any substance that undergoes a chemical reaction is called a _____________________. A substance that is formed is called a _____________________. Chemists use chemical equations to communicate what is happening in a chemical reaction. Chemical equations come in several forms. Word Equations
A word equation identifies the reactants and products by name. For example, sodium + chlorine → sodium chloride In this equation, “+” means “reacts with” and “→” means “to form”. Practice Problems
1. Describe each reaction using a word equation. Label the reactants and products.
a. calcium and fluorine react to form calcium fluoride
b. barium chloride and hydrogen sulphate react to form hydrogen chloride and barium sulphate
c. calcium carbonate, carbon dioxide and water react to form calcium hydrogen carbonate
d. hydrogen peroxide reacts to form water and oxygen
e. sulfur dioxide and oxygen react to form sulfur trioxide
NOTES
Word equations are useful because they identify the products and reactants in a chemical reaction. They do not, however, provide any chemical information about the compound or elements themselves. Another shortcoming of the word equation is that some names for chemicals are very long and awkward to work with. Skeletal Equations A skeletal equation lists the chemical formula of each reactant on the left, separated by a “+” sign if more than one reactant is involved, followed by an arrow “→”. The chemical formula for each product is listed on the right, again separated by a “+” sign if more than one product is formed. A skeletal equation shows the state of each reactant by using the appropriate subscript.
For example, the reaction of sodium metal with chlorine gas to form sodium chloride can be represented by the following skeletal equation:
Na (s) + Cl 2 (g) → NaCl (s)
A skeletal equation is more useful to a chemist than a word equation because it show the formulas of the compounds involved. It also shows the state of each substance.
SYMBOL
+
→
(s)
(l)
(g)
(aq)
MEANING
Practice Problems
20. Write a skeletal equation for each of the following:
a. solid zinc reacts with chlorine gas to form solid zinc chloride
b. solid calcium and liquid water react to form solid calcium hydroxide and hydrogen gas.
c. Solid barium reacts with solid sulfur to produce solid barium sulfide
d. Aqueous lead (II) nitrate and solid magnesium react to form aqueous magnesium nitrate and solid lead.
21. In each reaction below, a solid reacts with a gas to form a solid. Write a skeletal equation for each reaction.
a. carbon dioxide + oxygen → calcium carbonate
b. aluminum + oxygen → aluminum (III) oxide
c. magnesium + oxygen → magnesium oxide
NOTES
NOTES
Why Skeletal Equations Are Incomplete Although skeletal equations are useful, they do not fully describe chemical reactions. For example, according to the skeletal equation showing the formation of sodium chloride, one molecule of sodium reacts with one chlorine molecule containing two chlorine atoms. The product is one formula unit of sodium chloride containing one atom of sodium and one atom of chlorine. Where has the extra chlorine atom gone ?
The Law of Conservation of Mass All atoms must be accounted for according to the Law of Conservation of Mass. The Law of Conservation of Mass states that in any chemical reaction, the_____________________
______________________________________________________________________________. In other words, according to this law, matter cannot be _____________________ nor _____________________. Balanced Chemical Equations A balanced chemical equation reflects the Law of Conservation of Mass. This type of equation shows that there is the same number of each type of atom on both sides of the equation. Some skeletal equations are already balanced. For example, the skeletal equation for the reaction between carbon and oxygen to form carbon dioxide shows one carbon and two oxygen atoms on the reactant side of the equation, and one carbon and two oxygen atoms on the product side of the equation
Fig 4.1
p. 114 text
Most chemical equations, however, are not balanced, such as the one showing the formation of sodium chloride.
Fig. 4.2, 4.3, 4.4
p.p. 115 of text delete writing in between
You cannot balance an equation by changing any of the chemical formulas. The only way to balance an equation is to put the appropriate numerical coefficient in front of each compound or element in the equation.
Practice Problems
22. Copy the following skeleton equations and balance them.
a. S (s) + O 2 (g) → SO 2 (g)
b. P 4 (s) + O 2 (g) → P 4 O 10 (s)
c. H2 (g) + Cl 2 (g) → HCl (g)
d. SO 2 (g) + H2 O (l) → H2 SO4 (aq)
23. Indicate whether these equations are balanced. If they are not, balance them.
a. 4 Fe(s) + 3 O 2 (g) → 2Fe2O3(s)
b. HgO(s) → Hg (l) + O 2 (g)
c. H2 O2 → 2 H2 O (l) + O 2 (g)
d. 2 HCl (aq) + Na2SO3 (aq) → 2 NaCl (aq) + H2 O (l) + SO 2 (g)
NOTES
Steps for Balancing Equations
Here are some steps to follow when balancing equations that are more complex:
1. Balance the element, other than hydrogen and oxygen that has the greatest number of atoms in reactant or product.
2. Balance the other elements other than hydrogen and oxygen
3. Balance hydrogen and oxygen, whichever is present in the combined state. Leave until last, whichever one is in uncombined state
4. Check that each equation is balanced by counting the number of atoms of each element on each side of the equation
When the equation is balanced, the coefficients should be in their lowest terms.
Some balancing equations from SCH3a workbook 2000
p. 117 Text
p. 118 Text
4.2
SYNTHESIS AND DECOMPOSITION REACTIONS
(p.p. 119 – 125 Text) Just as there are different types of compounds, there are different types of chemical reactions. We will learn about ______ major classifications used for chemical reactions. Synthesis Reactions In a synthesis reaction, two or more elements or compounds combine to form a new substance. Synthesis reactions are also known as combination or formation reactions. A general equation for a synthesis reaction is In a simple synthesis reaction, one element reacts with one or more other elements to form a compound. There are several types of synthesis reactions and recognizing the patterns of the various types will help you predict whether substances will take part in a synthesis reaction. When a metal or a non-metal react with oxygen, the product is an _____________________. For example, the rusting of iron reacts according to the following equation:
A second type of synthesis reaction involves the reaction of a metal and a non-metal to form a binary compound. For example,
2 K (s) + Cl2 (g) → 2 KCl (s)
There are many types of synthesis reactions in which one or more compounds are the reactants. We will look at synthesis reactions of this type involving compounds oxides and water. When a non-metallic oxide reacts with water, the product is an _________________. The acids that form when a non-metallic oxide and water react are composed of hydrogen cations, and polyatomic anions containing oxygen and a non-metal. For example, one contributor to acid rain is sulfuric acid that is formed when sulfur trioxide reacts with water:
SO3 (g) + H2O(l) → H 2 SO4 (aq)
When a metallic oxide reacts with water, the product is a metal hydroxide, which is a ______________. For example, in the formation of calcium hydroxide from calcium oxide:
CaO (s) + 2 H2O (l) → Ca(OH) (2) (aq)
Practice Problems p. 122 Text
Notes
Notes
Decomposition Reactions
In a decomposition reaction, a compound breaks down into elements or other compounds. Therefore, a decomposition reaction is the ________________________of a synthesis reaction. The general formula for a decomposition reaction is:
_______________________________________
The substances produced by a decomposition reaction can be either elements or other compounds. For example:
2 H2O → 2 H2 + O2 or the explosive decomposition of ammonium nitrate:
NH4NO3 (s) → N2O(g) + 2 H2O (g)
Practice Problems p. 123
Notes
Notes
Combustion Reactions A complete combustion reaction is the reaction of a compound or element with _____ to form the most common oxides of the elements that make up the compound. Combustion reactions are usually accompanied by the production of light and heat. The complete combustion of any compound containing carbon, hydrogen, and oxygen (such as ethanol, C2H2OH) produces ________________________________and ________________________. Compounds that contain elements other than carbon also undergo complete combustion to form stable oxides. In the absence of sufficient oxygen, carbon containing compounds undergo incomplete combustion leading to the formation of carbon ________________________, ________, and water. Carbon monoxide is a deadly gas. You should always be sure that you have sufficient oxygen in your indoor environment for your gas furnace, gas stove, or fireplace.
Practice Problems p. 124
Notes
Notes
Section Review
p. 125 Text
Notes
Notes
4.3
SINGLE DISPLACEMENT AND DOUBLE DISPLACEMENT REACTIONS
(p.p. 126 – 141 Text) Single Displacement Reactioons In a single displacement reaction, one element in a compound is ________________________ (or replaced) by another element. Two general reactions represent two different types of single displacement reactions. One involves a metal replacing a metal cation in a compound; For example:
Or an example
The second type of single displacement reaction involves a non-metal, usually a halogen, replacing an anion in a compound as follows:
Or an example
Single Displacement Reactions and the Metal Activity Series
Single displacement reactions 1, ,2 3
p. 126 text When analyzing single displacement reactions, use the following guidelines:
• Treat hydrogen as a metal
• Treat acids such as HCl as ionic compounds of the form H + Cl or sulfuric acid
(H2SO4 as H + H + SO42- )
• Treat water as ionic with the formula H + (OH - )
Practice Problems p. 127
Notes
Through experiments, chemists have ranked the relative reactivity of metals, including hydrogen in acids and water), in an activity series. The reactive metals, such as potassium, are at the top of the activity series while the unreactive metals such as gold are at the bottom.
The Metal Activity Series The more reactive metals are at the ________________________of the activity series, the more stable ones are at the bottom. A reactive metal will displace any metal in a compound that is ________________________it in the activity series. Metals from lithium to sodium will replace hydrogen as a gas from water. Metals from magnesium to lead will displace hydrogen as a gas only form acids. Copper, mercury, silver, and gold will not displace hydrogen from acids. A single displacement reaction always favors the production of the ____________ reactive metal. In other words, the free metal that is formed from the compound must always be less reactive than the metal that displaced it. For example:
2 AgNO3 (aq) + Cu (s) → Cu(NO3)2 (aq) + 2 Ag (s)
Silver metal is more stable than copper metal. In other words, silver is below copper in the activity series.
Table 4.2
p. 130 text
Inv 4-A
p.p. 128
Inv 4-A
p.p. 129
You can use the activity series to help you predict the products of a reaction of a metal and a metal containing compound. For example,
Fe(s) + CuSO4 (aq) →
Iron is above copper in the activity series, meaning iron can replace copper and the reaction will proceed.
Fe(s) + CuSO4 (aq) → FeSO4 (aq) + Cu(s)
In the following reaction involving silver and calciu chloride,
Ag (s) + CaCl2 (aq) →
Silver is below calcium on the activity series, meaning that it is less reactive, therefore no reaction would take place.
Practice Problems
Top of p.131 text
NOTES
Single Displacement Reactions Involving Halogens Non-metals, typically halogens, can also take part in single displacement reactions. For example, molecular chlorine can replace bromine from KBr, an ionic compound, producing bromine and potassium chloride.
Cl2(g) + 2 KBr(aq) → 2 KCl(aq) + Br2(l)
The activity series for the halogens directly ________________________the position of the halogens in the periodic table. It can also be shown in the following way. ________________________ is the most reactive, and ________________________is the least reactive.
F > Cl > Br > I
The activity series for halogen can be used the same way the activity series for metals is used. If the element is above another in the activity series, it will replace it. For example,
F2 (g) + 2NaCl (aq) → 2NaF (aq) + Cl2 (aq)
If the element is below another in the activity series, it will not replace it and no reaction will occur. For example,
I2 + CaBr2 (aq) → no reaction
Practice problems
Bottom of
p. 131 text
NOTES
Notes
Double Displacement Reactions
A double displacement reaction involves the ________________________of cations between two ionic compounds, usually in aqueous (water) solution. A general equation for a double displacement reaction is:
____________________________________________
If you have 2 unlabelled beakers, one with distilled water, the other with salt water. You could determine the salt water solution by adding a few drops of silver nitrate solutions. The formation of a white precipitate (solid) indicates the presence of sodium chloride. The following double displacement reaction has occurred:
Fig . 4.15
p. 132 text
Since silver chloride is virtually insoluble in water, it forms a solid compound or precipitate. Double displacement reactions tend to occur in aqueous solutions. You can tell a double displacement reaction has taken place in the following cases:
• A solid (precipitate) forms
• A gas is produced
• Some double displacement reactions form a molecular compound like water. It is hard to tell when water has been formed, because often the reactions take place in water.
Double Displacement Reactions that Form a Precipitate A precipitate is a solid that separates from a solution as a result of a chemical reaction. Many double displacement reactions involve the formation of a precipitate. To predict whether a precipitate will form in a double displacement reaction, you must know the solubilities of different compounds. To determine the products, and their physical states in a double displacement reaction, you must first “deconstruct” the reactants, and then “reconstruct” the cations by switching the cations. You will be given the information on whether any of the products will form a precipitate.
Given the following reactants, how would you predict the products of the reaction and their state ? (Note, many hydroxide compounds are insoluble. Potassium cations form soluble substances with all anions)
MgCl2(aq) + KOH(aq) →
Fig. 4.16
p. 133 text If both products are soluble, and neither product precipitates out, both ionic compounds are dissolved in water and no reaction occurs.
Practice problems
p. 134 text
Notes
NOTES
Double Displacement Reactions that Produce a Gas In certain cases, you know that a double displacement reaction has occurred because a gas is produced. The gas is formed when one of the products of the double displacement reaction decomposes to give water and a gas. For example, the reaction between sodium carbonate and hydrochloric acid. If you carry out this reaction, you immediately see the formation of carbon dioxide gas. The first reaction that takes place is a double displacement reaction. Determine the products in the following way.
• Separate the reaction into ions, and switch the anions. Write the chemical formulas for the products and balance the equation.
Na2CO3 (aq) + 2HCl (aq) → 2NaCl (aq) + H2CO3 (aq)
The carbonic acid is unstable and decomposes to carbon dioxide and water.
H2CO3 (aq) → H2O (l) + CO2 (g
Practice Problems
p. 135 text top of page
Notes
The Formation of Water In Neutralization Neutralization reactions are a special type of double displacement reaction that produces water. Neutralization involves the reaction of an acid with a base to form water and an ionic compound. For example, the neutralization of nitric acid with sodium hydroxide (a base) is a double displacement reaction.
HNO3 (aq) + NaOH (aq) → NaNO3 (aq) + H20 (l)
Often, neutralization reactions produce no precipitate or gas.
practice problems bottom p. 135 text
Notes
Investigation 4-B p. 146 text
Inv 4-B p. 137 text
Inv. 4-c p. 138 Tetx
Inv 4-v p. 139 text
Concept Organizer
p. 140 text Section Review Questions
p. 140 text
Section Review Questions
p. 141 text
Notes
Notes
Chapter 4 Review p. 149 text
Ch 4 Review
p. 150
NOTES
NOTES
Unit 1 review
p. 154 text
Unit 1 review
p. 155 text
Unit 1 review
p. 156 text
Unit 1 review
p. 157 text
Notes
Notes
Notes
Notes
Date: Name: Class:
Problem Solving STSE Connections in the Great Lakes
Chapter 1
BLM 1-1
Goal
Outline the connections among science, technology, society, and the environment in relation to the Great Lakes.
Procedure
1. Choose an Area of Concern or a more general issue from those listed below. Use print and electronic resources to investigate the chemicals involved, and the connections between science, technology, society, and the environment.
Great Lakes Areas of Concern
• Hazardous waste sites along the St. Lawrence river in Cornwall, Ontario, and in the Port Hope harbour, remain from past industrial waste handling practices.
• Sewage treatment plants and sewage overflow contributes to the pollution in the Bay of Quinte.
• Copper mining, milling, and smelting operations have contaminated the water in Torch Lake, connected to Lake Superior.
• Industrial chemical discharges pollute the Niagara River connecting Lake Erie and Lake Ontario.
• Urban run-off from Metro Toronto contaminates the water in Lake Ontario.
General Pollution Issues involving the Great Lakes
• The atmospheric/airborne deposition of toxic chemicals may have a large role in the pollution of the Great Lakes.
• Biomagnification and bioaccumulation of toxic chemicals through the food chain puts predators such as lake trout, salmon, herring gulls, and fish-eating humans at risk. (Hint: Check out fish consumption advisories for the Great Lakes.)
• Nuclear power plants discharge heated water into the Great Lakes. How does this affect the environment? Where is this happening?
2. Organize and present your findings using one or more media such as an essay, a web site, a PowerpointTM or CorelTM presentation, or a brochure.
Analysis
1. What chemicals are involved in the issue you investigated?
2. For what purpose were the chemicals originally produced?
3. How did the chemicals benefit society?
4. How did the chemicals affect the environment?
5. What is being done to solve the problem?
Conclusions
6. Use a concept web to illustrate the STSE connections present in this issue.
UNIT 2
CHEMICAL
REACTIONS
UNIT 1
MATTER AND CHEMICAL BONDING
UNIT 1
MATTER AND CHEMICAL BONDING
1.1 The Study of Chemistry
What is chemistry ?
1.2
Define matter:
What are properties?
What is a physical property?
A physical property is s property that you can observe without changing one kind of matter into something new.
What is a chemical property?
Using Measurement to Describe Matter
What is the name of the system of measurement that scientists rely on to communicate effectively? What does it allow scientists to do?
Name five of the base S.I. units?
Complete the table below:
Table 1: Important SI Quantities and Their Units Quantity Definition S.I. Units or their derived equivalents Equipment used to measure the quantity
mass
length
temperature
volume
Mole
Density
Energy
Measurement and Uncertainty
What are exact numbers?
What two factors affect your ability to communicate measurements and calculations?
Significant Digits, Certainty, and Measurements
Calculated answers have the same number of significant digits as the least-certain measurement involved in the calculation; digits in excess of the number of least-certain digits must be rounded off when expressing the final answer.
The digits that you record when you measure something are called significant digits. Significant digits include all the digits that you are certain about and a final, uncertain digit that you estimate.
How Can You Tell which Digits are Significant ?
Complete the following table of rules for determining significant digits:
Rules Examples
Explaining Three Significant Digits in the Example Above
Explaining Five Significant Digits in the Example Above
Complete Practice Problems #1, 2 (text p. 18)
What is Precision?
Complete Investigation 1-A Observing Aluminum Foil (text p. 13)
Calculating With Significant Digits (text p. 20 – 22)
In this course, you will often take measurements and use them to calculate other quantities. You must be careful to keep track of which digits are significant.
Why is it important to keep track of significant digits?
Complete the rules for reporting significant digits in calculations in the table below:
Table 1 Rules for Reporting Significant Digits in Calculations
Rule 1 Multiplying and Dividing
Rule 2 Adding and Subtracting
Rule 3 Rounding
Sample Problem Reporting Volume Using Significant Digits
A student measured a regularly shaped sample o f iron and found it to be 6.78 cm long, 3.906 cm wide and 11 cm tall. Determine its volume to the correct number of significant digits.
Length = 6.78 cm (3 sig digits)
Width = 3.906 cm (4 sig digits)
Height = 11 cm (2 sig digits)
Solution
Sample Problem Reporting Mass Using Significant Digits
Suppose that you measure the masses of four objects as 12.5 g, 145.67 g, 79.0 g, and 38.438 g.
What is the total mass of the four objects ?
Strategy:
Complete Practice Problems # 3 a – g, p. 22 text) Reinforcement Reporting Significant Digits in Calculations
Chapter 1
BLM 1-2
Goal
Reinforce your understanding of significant digits.
Procedure
1. State the number of significant figures in each of the following:
(a) 3570
(b) 17.505
(c) 41.400
(d) 0.51
(e) 0.000 572
(f) 0.009 00
(g) 41.50 × 10-4
(h) 0.007 160 × 105
(i) 1.234 00 × 108
(j) 0.000 410 0 × 107
2. Perform the following operations and give the answer to the correct number of significant digits.
(a) 15.1 + 75.32
(b) 178.904 56 – 125.805
(c) 4.55 × 10-5 – 3.1 x 10-5
(d) 0.000 159 + 4.0074
(e) 1.805 × 104 + 5.89 × 104
(f) 0.000 817 - 0.000 048 1
(g) 8.166 × 105 – 7.819 × 105
(h) 45.128 + 8.501 87 – 42.18
(i) 5.677 × 10-6 + 7.785 × 10-6
(j) 8.75 × 10-9 + 6.1157 × 10-9
(k) 1.99 ÷ 3.1
(l) 1200.0 ÷ 3.0
(m) 5.32 x 10-4 ÷ 4.218 × 10-8
(n) 45.32 x 2.3
(o) 0.024 00 ÷ 6.000
(p) 12.4 x 0.30
(q) (5.50 x 108) × (4 x 105)
(r) 7.4 ÷ 3
(s) 4.75 × 5
(t) 2.5 × 6.700 ÷ 0.891
Date: Name: Class:
Assessment
Chapter 1
BLM 1-4
Goal
Demonstrate and assess your understanding of the concepts presented in Chapter 1
Part 1:
Multiple
Choice
Circle the letter corresponding to the best answer.
1. The metric base unit to measure linear dimensions is the:
(a) kelvin
(b) metre
(c) candela
(d) kilogram
(e) second
2. Which of the following is not an SI base unit?
(a) litre
(b) second
(c) ampere
(d) kilogram
(e) mol
3. Which digit or digits are uncertain in the result of a scientific measurement?
(a) the last
(b) the last two
(c) all of them
(d) none of them
(e) the first
4. How many significant digits are contained in the number 0.000540?
(a) two
(b) three
(c) five
(d) six
(e) seven
5. Which of the following is characteristic of pure substances?
(a) They are solutions.
(b) They are mixtures.
(c) They are composed of one element.
(d) They have a fixed chemical composition.
(e) They have variable composition.
6. All mixtures:
(a) are easy to separate
(b) have a constant boiling point
(c) are composed of parts with different identities
(d) have a chemical formula
(e) are the result of chemical change
7. An example of a pure substance is:
(a) salt water
(b) copper(II) chloride
(c) air
(d) tap water
(e) soda water
Part 2: Short Answer & Calculations 8. Carry out the following metric conversions.
(a) 45.6 g to mg
(b) 0.00867 kL to mL
9. Perform the following calculations and express the answer to the correct number of significant digits.
(a) 0.000354 × 2.8
(b) 7.90 + 3.222
(c) 8.432 ÷ 2.5
(d) 0.17 – 0.154
10. Identify the following as being either a physical or a chemical change.
(a) a chicken cooking in the oven
(b) ice cubes subliming
(c) iron rusting in air
(d) water evaporating form salt water in a beaker when heated
(e) toasting a slice of bread
(f) burning a log
11. Round off the following to the number of significant digits indicated.
(a) 0.0002554, to 2 significant digits
(b) 35.8348, to 4 significant digits
(c) 11.4555, to 3 significant digits
12. How long is the object shown below? Record your answer using the correct number of significant digits.
1.3 Classifying Matter and Its Changes Matter changes in response to changes in energy.
What is energy?
What is a physical change?
What is a chemical change
Classifying Matter
What is a mixture ?
What is a heterogeneous mixture?
What is a homogeneous mixture? What’s another name given?
What is a pure substance?
How are pure substance further classified?
What is an element? Give some examples.
What is a compound? Give some examples.
Chapter 1 Review (text p.p. 29 – 31) #2, 5, 6, 7, 8, 9, 10, 11, 12,
Chapter 2 Elements of the Periodic Table (text p. 33 – 64)
2.1 ATOMS AND THEIR COMPOSITION (text p.p. 34 – 39)
THE MODERN VIEW OF THE ATOM
What is the smallest unit of an element that still retains the properties of that element?
Atoms themselves are made of even smaller particles. Name the three subatomic particles and give their location in the atom.
THE ATOMIC THEORY OF MATTER
Who was John Dalton and what is he famous for? What was his model called?
Describe the four main points of Dalton’s atomic theory.
Dalton’s concept of the atom was one of several ideas in his atomic theory of matter. Dalton’s atomic theory of matter (1809) has four main points:
•
• •
•
• Law of Conservation of Mass:
• Law of Definite Proportions:
Complete the table below:
Table 2.1 Properties of Protons, Neutrons, and Electrons
Subatomic Particle
Charge
Symbol
Mass (g)
Radius (m)
Electron
Proton
Neutron
EXPRESSING THE MASS OF SUBATOMIC PARTICLES • What is an Atomic Mass Unit
The Nucleus of an Atom
• What does Atomic Number represents? ( Z )
• What does the Mass Number represents? ( A )
Information about an element’s protons and neutrons is often summarized using the chemical notation shown in Figure 2.3.
Mass # A
X
Atomic # Z
In the figure above, which letter represents the atomic symbol of the element ?
In the following example, identify the mass number and the atomic number
Atomic number
Mass number
19 F 9
For the isotope of fluorine above, what do the numbers represent?
How can you determine the number of neutrons in the nucleus of an atom?
How many neutrons would this isotope of fluorine have? PRACTICE PROBLEMS
1. Complete the missing values or chemical notations about the atoms from the table on p. 37 of your textbook below.
a. b. c.
d. e f.
g h i
j k l
m n o
p. q. r.
s. t. u.
v. w. x.
When chemists refer symbolically to an element, they will often leave out the atomic number. For example, for oxygen, they would write 16 O or O – 16.
USING THE ATOMIC NUMBER TO INFER THE NUMBER OF ELECTRONS
How can you infer the number of electrons from the atomic number? Why can you do this?
ISOTOPES AND ATOMIC MASS
What is am isotope?
Compare the chemical properties of isotopes of the same element. Explain why it is like this.
What is radioactivity”?
Atomic Weights
What information does the average atomic weight of an element give? What other name can it have?
Write the formula used to calculate the average atomic weight.
Example 1 Sulfur has three isotopes which have the following relative abundances: S – 32 = 95% S – 33 = 0.8 % S – 34 = 4.2 % Calculate the atomic weight of sulfur.
ELECTRONS IN ATOMS
What subatomic particle is responsible for the chemical properties of an element?
REVISITNG THE ATOMIC THEORY
John Dalton did not even know about subatomic particles when he developed his atomic theory. Even so, the modern atomic theory, shown below, retains many of Dalton’s ideas, with only a few modifications.
The Modern Atomic Theory
SECTION REVIEW
1. For the table below, fill in the missing information in the appropriate space provided.
ELEMENT ATOMIC NUNMBER
MASS NUMBER
NUMBER OF PROTONS
NUMBER
OF ELECTRONS
NUMBER OF NEUTRONS
(a)
(b)
108
(c)
47
(d)
(e)
(f)
(g)
33
(h)
42
(i)
35
(j)
(k)
(l)
45
(m)
79
179
(n)
(o)
(p)
(q)
(r)
(s)
(t)
50
69
1. a. b. c. d.
e f. g. h.
i. j. k. l.
m. n. o. p.
q. r. s. t.
2. Explain the difference between an isotope and a radioisotope. Provide an example other than oxygen to support your answer.
3. Examine the information presented in the following pairs
3 3 14 16 19 18 H and He C and N F and F 1 2 6 7 9 9
(a) For each pair, do both members have the same number of protons ? electrons ? neutrons?
(b) Which pair or pairs consist of atoms that have the same value for Z? Which consists of atoms that have the same value for A ?
4. Compare Dalton’s atomic theory with the modern atomic theory. Explain why scientists modified Dalton’s theory.
5. Fill in the following table:
Protons
Electrons
Neutrons
P – 31
Al – 27
H-1 (protium)
H – 2 (deuterium)
EXPANDING THE MODEL OF THE ATOM (p.p. 656 – 659)
Emission Spectra
What can each wavelength of visible light be re[resented by?
What is a bright-line spectrum ?
Bohr’s Model of the Atom In 1913, Danish physicist, Niels Bohr developed a model of the atom that explained hydrogen emission spectrum.
Describe the location and motion of electros in Bohr’s a toms
The following three points of Bohr’s theory help explain hydrogen’s emission spectrum:
• In Bohr’s model, atoms have specific allowable ______levels. He called these energy levels stationary states. Each of these levels corresponds to a fixed, circular orbit around the nucleus.
• An atom does not give off energy while its electrons are in a stationary state
• An atom changes states by giving off, or absorbing a quantity of light energy exactly equal to the ______in energy between the two stationary states.
Bohr’s model was revolutionary because he proposed that the energy absorbed or emitted by an atom needed to have ______ ______. The energy change was ______, rather than continuous. When something is quantized, it is limited to discrete amounts or multiples of discrete amounts. Two great scientists paved the way for this surprising idea. German physicist Max Plank had already proposed that light was quantized, meaning that it exists in packets. Building on this idea, Einstein proposed that light could behave as particles, which he called photons. The energy associated with the light in a line spectrum corresponds with the change in energy of an electron as it moves _______or _______________ an energy level. For example, when electrons in hydrogen atoms that have been excited to the third energy level and subsequently drop to the second energy level, they emit light that has a specific wavelength. They emit photons of _____________ light that have a wavelength of 656.5 nm. These photons cause the red line spectrum for hydrogen seen in Figure D1 above.
Why was a new model required?
Sublevels The spectra of many electron atoms suggested that a more complex structure was needed. Notice that in Figure D1, the spectra for these more complex atoms have groups of lines close together. The groups of lines are separated by spaces. The large spaces represent energy differences between ______ ______, while the smaller spaces represent energy ______within levels. If the electrons are changing energy within the levels, this suggests that there are sublevels within each level., each with its own slightly different energy. The idea of different energy levels 1, 2, 3, 4, . . . remains, but each energy level is split up into sublevels called _____, _____, _____, _____. Examine the table below to see how electrons are arranged in sublevels for energy levels 1 to 4.
Each sublevel has its own energy. Examine the figure on the next page to see the relative energies of the different levels and sublevels. But where would you find electrons in this new model? How are the sublevels oriented in space?
Particles with Wave Like Properties
In 1924, Louis de Broglie suggested that all matter had -like properties. De Broglie developed an equation that allowed him to calculate the wavelength associated with any object. Objects we can see and interact with have calculated wavelengths that are smaller than electrons. Their wavelengths are so tiny compared to their size, that they do not have any measurable effect on the motion of the objects
For very tiny moving particles such as electrons, however, the wavelength becomes very significant. De Broglie’s theory was proven by experiment when streams of electrons produced diffraction patterns similar to electromagnetic radiation that was already known to travel in waves.
Orbitals
In 1920, Erwin Schrodinger used de Broglie’s idea that matter has wavelike properties. Schrodinger proposed the quantum mechanical model of the atom. In this new model, he abandoned the notion of the electron as a small particle orbiting the nucleus. Instead, he described the behaviour of electrons in terms of wave functions.
This imprecise nature of Schrodinger’s model was supported shortly afterwards by principle proposed by Heisenburg. Heisneburg demonstrated that it is impossible to know both an electron’s path and its exact location. Therefore it is impossible to know the exact location of an electron at any time, without some degree of uncertainty. Heisenburg’s uncertainty principle shows that you can never know both the ______and the ______ of an electron beyond a certain measure of precision. Heisneburg also showed that if you could know either the velocity or position precisely, then the other property would be uncertain. Therefore, when talking about where electrons are found, you can only talk in terms of
______ ______.
Schrodinger used a mathematical wave equation to define the probability of finding an electron within an atom. There are multiple solutions to this wave equation, and Schrodinger called these solutions ______. Each solution provides information about the energy and location of an electron within an atom. Each orbital has a specific energy associated with it, and each contains information about where inside the atom, the electrons would spend most of their time.
These regions of space where an electron is most likely to be found are called ______. In solving the wave equation, three unknowns were encountered (a fourth was added later). These unknowns are called the ______ ______and describe different properties of the electron.
The Four Quantum Numbers
1 . The Principle Quantum Number (_____)
• Indicates the ______an electron has
• Tells us which energy ______the electron is in
• Can have values of n = ______
• The higher the ‘n’ value, the ______the energy
• Correspond to Bohr’s orbits
2. The ______Quantum Number (_____)
• Indicates the ______of the orbital
• It can have values of L = ______ ______
• n- value possible ‘L’ values
1
2
3
4.
Note: the total number of possible ‘L’ values is the same as the ‘n’ value L- VALUE
0
1
2
3
SHAPE
SPEHRICAL
DUMBBELL
DOUBLE DUMBBELL
??
??
NAME
____– orbital
____– orbital
____– orbital
____– orbital
____– orbital
3. The Magnetic Quantum Number (M L)
• indicates the orientation of the orbitals about the nucleus.
• The number of values indicates the number of orbitals that exist with a given
• M L can have the values of - L . . . 0 . . . + L
L
0
1
2
3
4
5
NAME
M L
NUMBER OF POSSIBLE ORBITALS
4. _________________Quantum Number (_____)
• Related to the _________________ of an electron
• All electrons spin on their axis. The two different spins are assigned the value of
________________________
• indicates that each orbital can hold a maximum of ________ electrons. If there are two electrons in an orbital, they must have opposite spins.
Complete the following summary table on the Quantum Theory below:
n
L
NAME
M L # of
Orbitals
M S # of
Electrons Total # of
Electrons
1
2
2
3
3
3
4
4
4
4
The ‘n’ Rules
• The nimber of different shapes (or ) of orbitals in a given energy level =
• The total number of orbitals in a given energy level =
• The total number of electrons in a given energy level =
Shapes of Orbitals The shapes of the probability graphs of Schrodinger’s wave functions are the shapes of the orbitals in which electrons reside in atoms. You can visualize orbitals as electron clouds. The shape of each cloud is based on probability – it tells where the electron spends most of its time. Examine the figure below. The ‘s’ orbitals are spherical in shape. Each of these spherical shells contains two electrons. There are three ‘p’ orbitals for each sublevel, each with a capacity for two electrons. Each ‘p’ orbital is shaped something like a dumbbell. For each subleve, the ‘p’ orbitals are oriented along the x, y, and z axes. The d’ and ‘f’ orbitals are quite complex in shape. Each ‘d’ sublevel contains 5 ‘d’ orbitals and each ‘f’ sublevel contains 7 ‘f’ orbitals.
Filling Orbitals
What is the maximum number of electrons that can occupy any one orbital?
State the Pauli Exclusion Formula
How do electrons fill orbitals within atoms? They do so to minimize the potential energy of the atom.
1. The maximum number of electrons that any orbital can hold is _____________
2. They will fill the ________________________energy level first. The 1s orbital will fill before the 2s orbital, which fills before the 2 p orbitals and so on.
1. When occupying two or more orbitals with the same energy (for example any of the 3 ‘p’ orbitals), electrons will half fill each orbital until all are half filled before adding a second electron to each one. This is called ________________________Rule. You can show how electrons fill orbitals using superscripts in a notation called the electron configuration of the atom. For example, a boron atom contains 5 electrons. The electron configuration for Boron would be 1s 2 2s 2 2p 1. The electron configuration for Nitrogen is 1s 2 2s 2 2p 3.
Electron Configurations
Write the electron configuration for the following elements:
1. 21 Sc
2. 23 V
3. 30 Zn
4. 35 Br
5. 50 Sn
6. 57 La
7. 56 Ce
8. 64 Gd
9. 50 Hg
10. 84 Po
Quantum Review 1
ATOMS, ELEMENTS, AND THE PERIODIC TABLE (Text p.p. 40 – 47) By the mid 1800’s, there were 65 known elements. Chemists studied these elements intensively and recorded detailed information about their reactivity and the masses of their atoms. Some chemists began to recognize patterns and behaviour of many of these elements.
What was the contribution of Dmitri Mendeleev to the science of chemistry?
What was one of the major differences between Mendeleev’d original periodic table ad the new one?
State the Periodic Law
PRACTICE PROBLEMS
2. Identify the name and symbol of the elements in the following location of the periodic table:
a. Group 14 (IVA), Period 2
b. Group 11 (IB), Period 4
c. Group 18 (VIIIA), Period 6
d. Group 1 (IA), Period 3
e. Group 12 (IIB), Period 5
f. Group 2 (IIA), Period 4
g. Group 17, (VIIA), Period 5
h. Group 13 (IIA), Period 3
ELECTRONS AND THE PERIODIC TABLE
We have seen how the periodic table organizes elements so that those with similar properties are in the same group. You have also seen how the periodic table shows a clear distinction among metals, non-metals, and metalloids. Other details of the organization of the periodic table may seem strange. Why, for example, are there different numbers of elements in the periods.
The reason for this, and other details of the periodic table’s organization, involves the number and arrangement of ________________________in atoms of each element.
ELECTRONS AND ENERGY LEVELS Electrons cannot move haphazardly. Their movement around an atomic nucleus is restricted to fixed regions of space. These regions of space are three-dimensional. The following figure shows a representation of these regions.
These regions are not solid, but rather are volumes of space in which electrons may be found. They are referred to as energy shells or energy levels. An electron that is moving in a lower energy level is found ________________________to the nucleus. It would have less energy than it would if it were moving in a higher energy level. There is a limit to the number of electrons that can occupy each energy level. For example, a maximum of _______________ electrons can occupy the first energy level. A maximum of ________________________electrons can occupy the second energy level. The periodic trends that result from organizing the elements by their atomic number are linked to the way in which electrons occupy and fill energy levels.
As shown in the figure above, a common way to show the arrangement of electron in atoms is to draw circles around the atomic symbol. Each circle represents an energy level. Dots represent electrons that occupy each energy level. This kind of diagram is called a Bohr –Rutherford diagram. Figure 2.9 B shows that the first energy level is full when _____________ electrons occupy it. Only two elements have two or fewer electrons: hydrogen and helium. Hydrogen has one electron, and helium has two. These elements, with their electrons in the first energy level, make up Period _______ of the periodic table. As you can see in Fig. 2.9 C, Period 2 elements have two energy levels. The second energy level is full when ________electrons occupy it. Neon, with a total of ten electrons, has its first and second energy levels filled. Notice how the second energy level fills with electrons as you move across the period from lithium to fluorine. Figure 2.9 C
PATTERNS BASED ON ENERGY LEVELS AND ELECTRON ARRANGEMENTS The structure of the periodic table is closely related to ________________________ levels and the arrangement of electrons. Two important patterns result from this relationship. One involves periods, and the other involves groups. The Period Related Pattern
Elements in Period 1 have electrons in how many energy levels?
Elements in Period 2 have electrons in how many energy levels?
How many energy levels could an element in Period 5 occupy?\
The Group Related Pattern The second pattern emerges when you consider the electron arrangements in the main- group elements: the elements in Group 1 (1A), 2 (2A) and 13 (3A) to 18 (8A). All the elements in each main group have the same number of electrons in their highest (outer) energy level. What is the name given to electrons that occupy the outermost energy levels of an atom?
What in the atom is responsible for the chemical behaviour of an element?
USING LEWIS STRUCTURES TO REPRESENT VALENCE ELECTRONS It is time consuming to draw electron arrangements using Bohr-Rughterford diagrams. It is much simpler to draw Lewis structures to represent elements and their valence electrons. To draw a Lewis structure, you replace the nucleus and the inner energy levels of an atom with its atomic symbol. Then you place dots around the atomic symbol to represent the valence electrons. The order you place the first four dots is up to you, however, place a dot on each side of the symbol before doubling up on any dot.
Practice Problems
2. On the blank periodic table, sketch the electron arrangements for the first 20 elements using Lewis structures.
3. Use the periodic table to draw Lewis structures for the following elements: barium (Ba), gallium (Ga), tin (Sn), bismuth (B), iodine (I), cesium (Cs), krypton (Kr) and xenon (Xe)
4. Identify the number of valence electrons in the outer energy levels of the following elements:
a. chlorine f. lead
b. helium g. antimony
c. indium h. selenium
d. strontium i. arsenic
e. rubidium j. xenon
THE SIGNIFICANCE OF A FULL OUTER ENERGY LEVEL Describe the electron arrangement of the noble gases. What property does this give to the nobe gases?
How many valence electrons to the noble gases have?
What is a stable octet?
Summary
You have seen that the structure of the periodic table is directly related to energy levels and arrangement of electrons. The patterns that emerge from this relationship enable you to predict the number of valence electrons for any main group element. They also enable you to predict the number of energy levels that an element’s electrons occupy.
Orbital Box Notation Using this method, each orbital is shown as a box with ________________________ representing the electrons. If there are two electrons in the same orbital box, then they must have opposite spins. This is shown by having the arrows facing in different _____________________.
Element Electron Configuration Orbital Box
4 Be
6 C
7 N
Hund’s Rule _______________________________________________________________________________________________________________________
8 O
9 F
21 Sc
50 Sn
What is the ground state electron configuration of an atom?
Excited State Configurations In order to obtain these configurations, energy must be added to the atom. When energy is added, one or more electrons is moved to a higher energy orbital. To which orbital the electron gets moved depends on the amount of energy that is added.
Write the ground state and a possible excited state electron configuration for the following isotope of magnesium.
12 Mg Ground State Electron Configuration Excited State (one of many)
Electron Configurations and Valences
Why do atoms react with other atoms?
What is it in the atom that determines its chemical properties?
How can an atom become more stable?
Describe several stable arrangements of electrons in atoms.
Stable Arrangements atoms include:
1.
2.
3.
4.
Valences Predict the most likely valences for the following:
1. 7 N
2. 16 S
3. 30 Zn
4. 50 Sn
Explain why:
i. Fe (# 26) has a (+2) and a (+3) valence. Which is more stable ?
ii. Rh (# 45) has a (+2) and a (+4) valence.
iii. Cerium (# 58) has a (+3 and a (+4) valence
Exceptions to the Predicted Electronic Configurations There are two families whose electron configurations are different than what would be predicted. They are the families VIB and IB. The elements in these families spontaneously rearrange their electrons to make themselves more stable. To do this, they promote an electron from an ‘s’ orbital and excite it to a ‘d’ orbital.
Examples
24 Cr
42 Mo29
29Cu
47 Ag
Section Review
1. State the Periodic Law and provide at least two examples to illustrate its meaning.
2. Identify the Group number for each of these sets of elements. Then write the symbols for the elements within it:
• alkali metals
• noble gases
• halogens
• alkaline earth metals
3. Identify the elements that is described by the following information. Refer to a periodic table as necessary.
• It is a Group 14 (IIIA) metalloid in the third period
• It is a Group 15 (VA) metalloid in the fifth period
• It is the other metalloid in Group 15
• It is a halogen that exists in the liquid state at room temperature
4. What is the relationship between electron arrangement and the organization of the elements of the periodic table ?
5 How many valence electrons are there in an atom of each of these elements ?
(a) neon sodium magnesium bromine chlorine silicon sulfur helium strontium tin
(b) Present your answers from Part A in the form of Lewis structures.
(c) Classify each element from Part A as a metal, non-metal, or a metalloid
6. Compare and contrast noble gases with other elements.
7.
(a) Draw Lewis Structures for each of the following elements: lithium potassium magnesium aluminum carbon
(b) Which of these elements have the same number of energy levels ?
(c) Which have the same number of valence electrons
8. Identify the elements with the following electron distributions by orbits:
a. 2, 8, 8 b. 2, 8, 6 c. 2, 4 d. 2, 8, 8, 2
5. What is meant by a stable octet ?
6. How many atoms must each of the following elements gain to achieve a stable octet ?
a. fluorine b. oxygen c. nitrogen d. carbon
7. How many electrons must the following atoms lose in order to achieve a stable octet
a. silicon b. phosphorus c. sodium d. argon
8. How does the sign of the valence relate to the gain or loss of electrons ?
9. Explain why atoms form ions.
2.3
PERIODIC TRENDS INVOLVING THE SIZES AND ENERGY LEVELS OF ATOMS In Section 2.1, we learned that the size of a typical atom is about 10 – 10 m, however, we also know that the atoms of each element are distinctly different. For example, the atoms of different elements have different numbers of protons. This means of course, that they also have different numbers of electrons. You may predict that the size of an atom is related to the number of protons and electrons it has. Is there any evidence to support this prediction, If so, is there a pattern that could help you to predict the relative size of an atom for any element on the periodic table ? Chemists define, and measure, an atom’s size in terms of its radius. The radius of an atom is the distance from its nucleus to the approximate outer boundary of the cloud-like region of its electrons. This boundary is approximate because atoms are not solid spheres. They do not have an outer boundary.
Trends for Atomic Size (Radius) There are two general trends for atomic size, describe them.
•
•
Practice Problems
7. Using only their location in the periodic table, rank the atoms in each set by decreasing atomic size. Explain your rankings. (a) Mg Be Ba
(b) Ca Se Ga
(c) Br Rb Kr
(d) Se Br Ca
(e) Ba Sr Cs
(f) Se Br Cl
(g) Mg Ca Li
(h) Sr Te Se
(i) In Br I
(j) S Se O
EXPERIMENT THE REACTIVITY OF ELEMENTS
Trends for Ionization Energy
What is ionization energy?
What is an ion?
Name two different types of ions formed by the atom. Explain how the two different types of ions are formed?
The figure below shows the formation of ions for several elements. As you examine the diagrams, pay close attention to
• the energy level from which the electrons are gained or given up
• the charge of the ion that is formed when an atom gains or gives up electrons
• the arrangement of the electrons that remain after the electrons are gained or given up. If you try to visualize the periodic table as a cylinder, (see Fig. 2.15 below) rather than a flat plane, you can begin to see the relationship between ion formation and the electron arrangement of the noble gases.
Describe how metals and non-metals of the main group tend to react. What are they trying to become like?
Give an example.
How do elements in Group 1A react?
How do elements in Group 1A react?
How is the charge on a cation related to its group number?
Describe a genera rule for how non-metals react.
The figure below shows the ionization energy that is required to remove one electron from the outer energy level of atoms of the main group elements. This energy is called the ____________-ionization energy. It is measured in units of kJ / mol. As you can see, atoms that give up electrons easily have _______________ ionization energies. You would probably predict that the alkali metals of Group 1 (1A) would have low ionization energies. These elements are, in fact, extremely reactive because it takes so little energy to remove their single valence electron.
All elements, except _________________________, have more than one electron that can be removed. Therefore, they have more than one ionization energy. The energy that is needed to remove a second electron is called the ________________________ionization energy. The energy needed to remove a third electron is called the ________________________ionization energy, and so on.
Summarizing Trends for Ionization Energy
Upon examination of the periodic trend for ionization energy it is the opposite of the trend for atomic radius.
Atomic Radius vs First Ionization Energy
Although there are a few exceptions, there are two general trends for ionization energy:
• Ionization energy tends to ________________________down a group as electrons in the outer energy level are further from the positive force of the nucleus, thus they are easier to remove than electrons in lower energy levels.
• Ionization energy tends to ________________________across a period. As you go across a period, the attraction between the nucleus and the electrons in the outer energy level increases. Thus, ______________________ energy is required to pull an electron away from its atom. For this trend to be true, you would expect the noble gases to have the highest ionization energies, and as is shown in Figure 2.16, they do.
Practice Problems
8. Using only a periodic table, rank the elements in each set by increasing ionization energy. Explain your answers. Explanation
(a) Xe He Ar
(b) Sn In Sb
(c) Sr Ca Ba
(d) Kr Br K
(e) K Ca Rb
(f) Kr Br Rb
9. Using only a periodic table, identify the atom in each of the following pairs with the lower first ionization energy.
(a) B O
(b) B In
(c) I F
(d) F N
(e) Ca K
(f) B Tl
Trends for Electron Affinity
Define electron affinity:
Electron affinity is the change in ________________________that occurs when an electron is added to the outer energy level of an atom to form a negative ion.
Figure 2.18 above gives the electron affinities of the main group of elements. If the ion that is formed by gaining an electron is stable, the electron affinity is expressed as a negative integer. The more ________________________ the ion, the higher is the negative integer for the electron affinity. Notice that fluorine has the highest electron affinity. This indicates that fluorine is very likely to be involved in chemical reactions. In fact, fluorine is the most ________________________of all elements. Metals have ___________ electron affinities. That is especially true for the Group 1 (1A) and 2 (2A) elements. Atoms of these elements form ________________________positive ions. A negative ion that is formed by elements of these groups is unstable. It breaks apart into a neutral atom and a free electron Periodic Trend for Electronegativity (p.p. 70 – 71)
Describe. the meaning of electro negativity
What symbol is used to represent electro negativity?
Compare the trend for electroonegativity with that of atomic radius .Summarise the trends.
Explain the trends for electro negativity and atomic radius>
Section Review
1. How does your knowledge of electron arrangement and forces in atoms help you explain the following periodic trends ?
a. atomic radius b. ionization energy c. electron affinity
2. Using only their position on the periodic table, rank each of the following sets of elements in order of increasing atomic size. Explain your answer in each case.
a. Mg S Cl b. Al B In
c. Ne Ar Xe d. Rb Xe Te
e. P Na F f. O S N
3. Using only their position on the periodic table, rank each of the following sets of elements in order of decreasing ionization energy. Explain your answer in each case.
a. Cl Br I b. Ga Ge Se
c. K Ca Kr d. Na Li Cs
e. S Cl Br f. Cl Ar K
4. Which element in each of the following pairs will have the lower electron affinity ? Explain your answer.
a. K or Cl b. O or Li
c. S or Se d. Cs or F
3.1
Classifying Chemical Compounds (p.p. 66 – 74) Most elements do not exist in nature in their pure form as ________________________. Gold, silver, and platinum are three metals that can be found in the earth’s crust as elements. They are called “precious metals” because their occurrence is so rare. Most other metals, and most other elements, are found in nature only as ________________________. There are only about 90 naturally occurring elements. In comparison, there are thousands upon thousands of different compounds in nature, and more are constantly being discovered. Because there are so many compounds, chemists have developed a classification system to organize them according to their properties, such as melting point, boiling point, hardness, conductivity and solubility.
Properties of Ionic and Covalent Compounds Based on their physical properties, compounds can be classified into two groups: ________________________compounds and ________________________compounds. Complete the properties of ionic and covalent compounds in the table below:
PROPERTY
IONIC COMPOUND
COVALENT COMPOOUNDS
State at room temperature
Melting Point
Electrical conductivity as a liquid
Solubility in water
Conducts electricity when dissolved in water
What is Bonding ?
What is a chemical bond?
Which electrons are involved in bonding?
Why do atoms form bonds?
What is responsible for the different properties of ionic and covalent bonds?
Describe the formation of an ionic bond.
Describe the formation of a covalent bond.
Predicting Bond Type Using Electronegativity (p. 72 Text)
How can you use electro negativity to predict bond type?
The Range of Electronegativity Differences
What is a pure covalent bond?
Describe the formation of an ionic bond.
Describe how chemist’s use electronegativity to classify bonds.
What is a polar covalent bond?
Practice Problems
1. Determine the ΔEN for each bond shown. Indicate whether each bond I ionic or covalent
a. O - H
b. C - H
c. Mg - Cl
d. B - F
e. Cr - O
f. C - N
g. Na - I
h. Na - Br
Section Review
1. Name the typical properties of an ionic compound. Give two examples of ionic compounds.
2. Name the typical properties of a covalent compound. Give two examples of covalent compounds.
3. Describe and explain the periodic trend for electronegativity.
4. Based only on their position in the periodic table, arrange the elements in each set in increasing attraction for electrons in a bond.
5. Determine the ΔEN for each bond. Indicate whether the bond is ionic or covalent.
a. N - O b. Mn - O c. H - Cl d. Ca - Cl
3.2 IONIC AND COVALENT BONDINNG: THE OCTET RULE
The Octet Rule When atoms form bonds, they are often more stable. We know that noble gases are the most stable elements in the periodic table. They are extremely non-reactive, and they tend not to form compounds. What the noble gases have in common is a _____________________ outer electron energy level. When an atom loses or gains electrons to achieve a filled outer electron energy level, the atom often becomes more stable. According to the octet rule, atoms bond in order to achieve an electron configuration that is the same as the electron configuration as a noble gas. When two atoms or ions have the same electron configuration, they are said to be isoelectronic. For example, Cl - is isoelctronic with Ar because both have 18 electrons and a filled outer energy level.
Define Isoelectronic:
Ionic Bonding Sodium has a very low electornegativity, and chlorine has a high electronegativty. Therefore, when sodium and chlorine interact, sodium transfers its valence electron to chlorine. As shown in the figure below, sodium becomes Na + 1 and chlorine becomes Cl -. When neutral sodium loses its one valence electron to chlorine, the resulting Na + 1 cation has an en electron energy level that contains ___________ electrons. It is _____________________ with the noble gas neon. On the other hand, chlorine’s outer energy level has seven electrons. When chlorine gains sodium’s electron, it becomes an anion _____________________ with the noble gas argon.
The figure below shows how to represent the formation of an ionic bond using Lewis Structures. Thus, in an ionic bond, electrons are _____________________ from one atom to another so that they form oppositely charged ions. The strong force of attraction between oppositely charged ions is what holds them together.
Practice Problems
1. For each bond below, determine the ΔEN. Is the bond ionic or covalent ? a. Ca - O
b. K - Cl
c. K - F
d. Li - F
e. Li - Br
f. Ba - O
2. Draw Lewis structures to represent the formation of each bond in Question 1
COVALENT BONDING
We have just learned what happens when the electronegativity difference between two atoms is greater than 1.7. The atom with the lower electronegativity transfers its valence electron(s) to the atom with the higher electronegativity. The resulting ions have opposite charges. They are held together by a strong ionic bond. Describe the bond formation between two atoms when the electronegativity difference is zero.
Describe how a coavalent bond is formed.
When two atoms of the same element form a bond, they share their electrons equally in a _____________________ covalent bond. Elements that bond to each other in this way are known as _____________________ molecules. When atoms such as carbon and hydrogen bond to each other, their electronegativities are so close that they share electrons almost equally. Carbon and hydrogen have an electronegativity difference of only 2.6 -= 2.2 = 0.4. In the figure below, you can see how one atom of hydrogen forms a covalent bond with four atoms of hydrogen. The compound methane, CH4 , is formed. Each hydrogen shares one of its electrons with carbon. The carbon shares its four electrons with each hydrogen atom. Thus, each hydrogen achieves a filled outer energy level, as does carbon. (Recall that elements in the first period only need two electrons to fill their outer energy level). When analyzing Lewis structures that show covalent bonds, count the shared electrons as if they belong to each of the bonding atoms.
Practice Problems
1. Show the formation of a covalent bond between two atoms of each diatomic element. a. iodine
b. bromine
c. hydrogen
d. fluorine 2. Use Lewis structures to show the simplest way in which each pair of elements forms a covalent bond, according to the octet rule. a. hydrogen and oxygen
b. chlorine and oxygen
c. carbon and hydrogen
d. iodine and hydrogen
e. nitrogen and hydrogen
f. hydrogen and rubidium
Mutliple Covalent Bonds (p. 82 Text) In covalent bonding, atoms sometimes need to share two or three pairs of electrons, according to the octet rule.
Describe how oxygen gas satisfies the octet rule.
Note: the number of bonds each atom must form is equal to its ________________ Section Review
1. Use Lewis structures to show how each pair of elements forms an ionic bond.
a. magnesium and fluorine b. potassium and bromine
c. rubidium and chlorine d. calcium and oxygen
2. Use Lewis structures to show how each pair of elements forms covalent bonds.
a. one silicon atom and two oxygen atoms
b. one carbon atom, one hydrogen atom, and three chlorine atoms
c. two nitrogen atoms
d. two carbon atoms bonded together with three hydrogen atoms bonded to one carbon atom and one hydrogen atom and one oxygen atom bonded to the other carbon.
3. Use what you know about electronegativity differences to decide what kind of bond would form between each pair of elements.
4. In general, the farther away two elements are form each other on the periodic table, the more likely they are to participate in ionic bonding. Do you agree with this statement? Explain why or why not.
Polar Covalent Bonds: The “In-Between Bonds”
When two atoms have an elecronegativity difference hat is greater than 0.5 but less than 1.7, they are considered to be a polar covalent bond. In a polar covalent bond, the atoms have significantly different electronegativities. The electronegativity difference is not great enough, however, for the less electronegative atom to transfer its valence electrons to the other, more electronegative atom. The difference is great enough for the bonding electron pair to spend more time near the more electronegative atom. For example, the bond between oxygen and hydrogen in water has an electronegativity difference of 1.24. Because this value is less than 1.7 but greater than 0.5, the bond is a polar covalent bond. The oxygen attracts the electrons more strongly than the hydrogen. Therefore, oxygen has a slightly negative charge and the hydrogen has a slightly positive charge. These charges are not full charges but rather a partial charge symbolized by the symbols δ + and δ -.
Practice Problems Predict whether each bond will be covalent, polar covalent, or ionic.
a. C – F
b. O – N
c. Cl – Cl
d. Cu – O
e. Si – H
f. Na – F
g. Fe – O
h. Mn – O
2. For each polar covalent bond in Question 1, indicate the location of the partial charges.
3. Arrange the bonds in each set in order of increasing polarity. (A completely polarized bond is an ionic bond)
a. H – Cl O – O Na – Cl
b. C – Cl Mg – Cl P – O N - N
3.4
Appendix 1 Periodicity and Bonding Questions
Unit 2
CHEMICAL
REACTIONS
CHAPTER 4: CLASSIFYING CHEMICAL REACTIONS: CHEMICALS IN BALANCE
4.1 Chemical Equations Any substance that undergoes a chemical reaction is called a _____________________. A substance that is formed is called a _____________________. Chemists use chemical equations to communicate what is happening in a chemical reaction. Chemical equations come in several forms. Word Equations
A word equation identifies the reactants and products by name. For example, sodium + chlorine → sodium chloride In this equation, “+” means “reacts with” and “→” means “to form”. Practice Problems
1. Describe each reaction using a word equation. Label the reactants and products.
a. calcium and fluorine react to form calcium fluoride
b. barium chloride and hydrogen sulphate react to form hydrogen chloride and barium sulphate
c. calcium carbonate, carbon dioxide and water react to form calcium hydrogen carbonate
d. hydrogen peroxide reacts to form water and oxygen
e. sulfur dioxide and oxygen react to form sulfur trioxide
2. What is one limitation of word equations?
Skeletal Equations Describe a skeletal equation.. Give an example.
A skeletal equation is more useful to a chemist than a word equation because it show the formulas of the compounds involved. It also shows the state of each substance.
Complete the table below that gives the meaning of the symbols in a chemical equation.
SYMBOL
+
→
(s)
(l)
(g)
(aq)
MEANING
Practice Problems
20. Write a skeletal equation for each of the following:
a. solid zinc reacts with chlorine gas to form solid zinc chloride
b. solid calcium and liquid water react to form solid calcium hydroxide and hydrogen gas.
c. Solid barium reacts with solid sulfur to produce solid barium sulfide
d. Aqueous lead (II) nitrate and solid magnesium react to form aqueous magnesium nitrate and solid lead.
21. In each reaction below, a solid reacts with a gas to form a solid. Write a skeletal equation for each reaction.
a. carbon dioxide + oxygen → calcium carbonate
b. aluminum + oxygen → aluminum (III) oxide
c. magnesium + oxygen → magnesium oxide
Why Skeletal Equations Are Incomplete Although skeletal equations are useful, they do not fully describe chemical reactions. For example, according to the skeletal equation showing the formation of sodium chloride, one molecule of sodium reacts with one chlorine molecule containing two chlorine atoms. The product is one formula unit of sodium chloride containing one atom of sodium and one atom of chlorine. Where has the extra chlorine atom gone?
State the Law of Conservation of Mass
Balanced Chemical Equations A balanced chemical equation reflects the Law of Conservation of Mass. This type of equation shows that there is the same number of each type of atom on both sides of the equation. Some skeletal equations are already balanced. For example, the skeletal equation for the reaction between carbon and oxygen to form carbon dioxide shows one carbon and two oxygen atoms on the reactant side of the equation, and one carbon and two oxygen atoms on the product side of the equation
Most chemical equations, however, are not balanced, such as the one showing the formation of sodium chloride. Na + Cl2 → NaCl
You cannot balance an equation by changing any of the chemical formulas. The only way to balance an equation is to put the appropriate numerical coefficient in front of each compound or element in the equation.
Practice Problems
22. Copy the following skeleton equations and balance them.
a. S (s) + O 2 (g) → SO 2 (g)
b. P 4 (s) + O 2 (g) → P 4 O 10 (s)
c. H2 (g) + Cl 2 (g) → HCl (g)
d. SO 2 (g) + H2 O (l) → H2 SO4 (aq)
23. Indicate whether these equations are balanced. If they are not, balance them.
a. 4 Fe(s) + 3 O 2 (g) → 2Fe2O3(s)
b. HgO(s) → Hg (l) + O 2 (g)
c. H2 O2 → 2 H2 O (l) + O 2 (g)
d. 2 HCl (aq) + Na2SO3 (aq) → 2 NaCl (aq) + H2 O (l) + SO 2 (g)
NOTES
Steps for Balancing Equations Here are some steps to follow when balancing equations that are more complex:
1. Balance the element, other than hydrogen and oxygen that has the greatest number of atoms in reactant or product.
2. Balance the other elements other than hydrogen and oxygen
3. Balance hydrogen and oxygen, whichever is present in the combined state. Leave until last, whichever one is in uncombined state
4. Check that each equation is balanced by counting the number of atoms of each element on each side of the equation
When the equation is balanced, the coefficients should be in their lowest terms.
Practice Problems (text p.p. 117 & 118) # 7, 8, 9
Section Review (text p. 118) # 2, 4
Synthesis and Decomposition Reactions
(text p.p. 119 – 125)
Just as there are different types of compounds are different types of chemical reactions. We will learn about five major classifications used for chemical reactions.
Synthesis Reactions
In a synthesis reaction, two or more elements or compounds combine to form a new substance.
Write the general equation for a synthesis reaction
In this course all you'll the with two specific types synthesis reactions involving compounds that you are familiar with -- oxides and water.
Describe the products formed when a non-metallic oxide reacts with water.
Describe the acids that are formed when a nonmetallic oxide and water react. Provide an example.
Describe the products formed when a metallic oxide reacts with water. Provide an example and write the balanced equation.
Practice Problems p. 122 # 10 – 13
Decomposition Reactions
Write the general formula for a decomposition reaction.
C → A + B
What is a decomposition reaction? Provide an example and a balanced equation
Practice Problems p.. 123 # 14 - 16
Combustion Reactions
Describe the products of a complete combustion reaction. Provide an example.
What is incomplete combustion?
Practice Problems p. 124 # 17- 20
Section Review P. 125 # 1 – 6
4.3
Single Displacement and Double Displacement Reactions
Describe a single displacement reaction.
Write the general reactions that represent two different types of single displacement reactions. Describe these two reactions.
Single Displacement Reactions and the Metal Activity Series
Most single displacement reactions involve one metal displacing another metal from a compound. In the following equation, magnesium metal replaces the zinc in ZnCl2, thereby liberating the zinc as the metal.
Mg(s) ZnCl2 (aq) → MgCl2 (aq) + Zn(s)
Writes three balanced equations that illustrate the various types of single displacement reactions involving metals. 1.
2.
3.
Write the general formula for a single displacement reaction
Write three guidelines that you should follow when analyzing single displacement reactions
•
•
•
Practice problems # 21 (text p. 127)
What is at activity series
Through experimentation chemists have ranked the we reactivity of the metals including hydrogen [in acids and water], in ant activity series.
Complete Investigation 4-A Creating an Activity Series of Metals
p.p. 128-129
The Metal Activity Series
Describe the arrangement of metals on the metal activity series
What does a single displacement reaction always favor?
Using the activity series predict the products of the following reactions:
Fe(s) + CuSO4 (aq) → FeSO4 (aq) + Cu(s)
Ag (s) + CaCl2(aq) →
Practice Problems (text p. 131 #22)
Single Displacement Reactions Involving a Halogens
Predict to the products for the following reaction:
Cl2 (g) + 2 KBr (aq) →
Describe the activity series for halogens.
Predict the products of the following reactions:
F2(g) + 2NaCl →
I2(g) + CaBr2(aq) →
Practice Problems: (text p. 131 #23, 24)
Single Displacement Reactions Involving Halogens Non-metals, typically halogens, can also take part in single displacement reactions. For example, molecular chlorine can replace bromine from KBr, an ionic compound, producing bromine and potassium chloride.
Cl2(g) + 2 KBr(aq) → 2 KCl(aq) + Br2(l)
The activity series for the halogens directly ________________________the position of the halogens in the periodic table. It can also be shown in the following way. ________________________ is the most reactive, and ________________________is the least reactive.
F > Cl > Br > I
The activity series for halogen can be used the same way the activity series for metals is used. If the element is above another in the activity series, it will replace it. For example,
F2 (g) + 2NaCl (aq) → 2NaF (aq) + Cl2 (aq)
If the element is below another in the activity series, it will not replace it and no reaction will occur. For example,
I2 + CaBr2 (aq) → no reaction
Practice Problems (Text p. 131)
Double Displacement Reactions
Describe what takes place in a double displacement reaction.
Write the general formula for a double displacement reaction.
An example of a double displacement reaction between sodium chloride and silver nitrate
Since silver chloride is virtually insoluble in water, it forms a solid compound or precipitate.
Double displacement reactions tend to occur in aqueous solutions. You can tell a double displacement reaction has taken place in the following cases:
What are three pieces of evidence that a double displacement reaction has occurred?
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• •
Double Displacement Reactions that Form a Precipitate
What is a precipitate ?
How can you predict whether a precipitate will be produced during a reaction?
Given the following reactants, how would you predict the products of the reaction and their state ? (Note, many hydroxide compounds are insoluble. Potassium cations form soluble substances with all anions)
MgCl2(aq) + KOH(aq) →
If both products are soluble, and neither product precipitates out, both ionic compounds are dissolved in water and no reaction occurs.
Practice Problems (text p. 134)
Double Displacement Reactions that Produce a Gas In certain cases, you know that a double displacement reaction has occurred because a gas is produced. The gas is formed when one of the products of the double displacement reaction decomposes to give water and a gas. For example, the reaction between sodium carbonate and hydrochloric acid. If you carry out this reaction, you immediately see the formation of carbon dioxide gas. The first reaction that takes place is a double displacement reaction. Determine the products in the following way.
• Separate the reaction into ions, and switch the anions. Write the chemical formulas for the products and balance the equation.
Na2CO3 (aq) + 2HCl (aq) → 2NaCl (aq) + H2CO3 (aq)
The carbonic acid is unstable and decomposes to carbon dioxide and water.
H2CO3 (aq) → H2O (l) + CO2 (g
Practice Problems (text p. 135 )
The Formation of Water In Neutralization
Neutralization reactions are a special type of double displacement reaction that produces water.
What is a neutralization reaction ? Give an example.
Often, neutralization reactions produce no precipitate or gas.
Practice Problems (text p. 135)
Investigation 4-B (text p.p. 136 – 137)
Investigation 4-c (text p. 138 )
Section Review Questions (text p. 140
Chapter 4 Review text p. 149
Unit 1 review (text p. 156)
Date: Name: Class:
STSE Connections in the Great Lakes
Goal
Outline the connections among science, technology, society, and the environment in relation to the Great Lakes.
Procedure
1. Choose an Area of Concern or a more general issue from those listed below. Use print and electronic resources to investigate the chemicals involved, and the connections between science, technology, society, and the environment.
Great Lakes Areas of Concern
• Hazardous waste sites along the St. Lawrence river in Cornwall, Ontario, and in the Port Hope harbour, remain from past industrial waste handling practices.
• Sewage treatment plants and sewage overflow contributes to the pollution in the Bay of Quinte.
• Copper mining, milling, and smelting operations have contaminated the water in Torch Lake, connected to Lake Superior.
• Industrial chemical discharges pollute the Niagara River connecting Lake Erie and Lake Ontario.
• Urban run-off from Metro Toronto contaminates the water in Lake Ontario.
General Pollution Issues involving the Great Lakes
• The atmospheric/airborne deposition of toxic chemicals may have a large role in the pollution of the Great Lakes.
• Biomagnification and bioaccumulation of toxic chemicals through the food chain puts predators such as lake trout, salmon, herring gulls, and fish-eating humans at risk. (Hint: Check out fish consumption advisories for the Great Lakes.)
• Nuclear power plants discharge heated water into the Great Lakes. How does this affect the environment? Where is this happening?
2. Organize and present your findings using one or more media such as an essay, a web site, a PowerpointTM or CorelTM presentation, or a brochure.
Analysis
1. What chemicals are involved in the issue you investigated?
2. For what purpose were the chemicals originally produced?
3. How did the chemicals benefit society?
4. How did the chemicals affect the environment?
5. What is being done to solve the problem?
Conclusions
6. Use a concept web to illustrate the STSE connections present in this issue.
UNIT C
QUANTITIES
IN
CHEMICAL
REACTIONS
5.1
Isotopes and Average Atomic Mass The mass of an atom is expressed in atomic mass units (u.). Atomic mass units are a relative measure, define by the mass of one carbon – 12 atom. One carbon – 12 atom has been assigned a mass of 12 u, or 1u = 1/12 the mass of one atom of carbon – 12. Average Atomic Mass
What is the average atomic mass of an element?
6 C 12.01
Atomic Weights
Provide the formula to calculate atomic weight. What do the symbols represent?
Sample Problem
Naturally occurring silver exists as two isotopes. From the mass of each isotope and the isotopic abundance listed below, calculate the average atomic mass of silver. ISOTOPE ATOMIC MASS (U) RELATIVE ABUNDANCE Ag-107 106.9 51.8 Ag – 109 108.9 48.2
Solution
Practice Problems p.167 & 168 # 1- 4
5.2
THE AVOGADRO CONSTANT AND THE MOLE
(p. 171 Text)
Mole (symbol mol). What number does the mole represent?
•
•
How many formulae units would one mole NaCl contain?
Practice Problems
11. If you drove for 6.022 x 10 23 days at a speed of 100 km/h, how far would you travel ?
12. If you spent $ 6.022 x 10 23 at a rate of $1.00 /s, how long, in years would the money last? Assume that every year has 365 days.
Converting Moles to Number of Particles Write the Formula Converting Moles to Number of Particle
Example Problem
A sample contains 1.25 mol of nitrogen dioxide NO2.
a. How many molecules are in the sample ?
b. How many atoms are in the sample ?
Solution
Practice Problems (text p. 177 )
Converting Number of Particles to Moles Write the Formula Converting Number of Particles to Moles
Sample Problem: How many moles are present in a sample of carbon dioxide, CO2, made up of 5.83 x 10 24 molecules? Solution
Practice Problems (text p. 178)
MOLES WORKSHEET
1. What is meant by the term of mole ?
2. What is meant by the term Molar Volume?
3. What Is meant by the term of Molar Mass?
4. How many atoms make up one molecule of the following?
a. NaCl b. CCl4 c. H3PO4
d. Fe2(SO4)3 e. Al2(Cr2O7) 3 f. Ca(OH) 2
5. Determine the Molar Mass each of the following
a. H2SO3 b. C3H8 c. Ca(HCO3) 2
d. MnSO4 e. U(CH3COO) 3
6.
a. How many molecules are present and 5.2 moles of H2O
b. How many atoms are present and 5.2 moles of H2O
c. How many molecules are present in three moles of NH3
d. How many activities are present in three moles of NH3
5.3
Molar Mass
One mole of an element has a mass expressed in grams numerically equivalent to the element’s average atomic mass expressed in atomic mass units. One mole of zinc atoms has a mass of 65.39 g. You can use the periodic table to determine the mass of one mole of an element. What is Molar Mass
Complete the table below
Element
Average Atomic Mass (u)
Molar Mass (g)
Hydrogen, H
Oxygen, O
Sodium, Na
Argon, Ar
The molar mass relates the amount of an element or compound in moles, to its mass.
Finding the Molar Mass of Compounds
Find the molar mass of BaO
Sample Problem
What is the mass of one mole of calcium phosphate, Ca3(PO4)2 ? Solution
Practice Problems p. 184 Text # 23 - 26
Investigation 5-A (text p.p. 182 & 183)
Complete Investigation 5-A and answer Analysis 1-6, Conclusion 7, and
Application 8, 9 in this workbook.
Converting from Moles to Mass
Write the formula that converts moles to mass.
Sample Problem
A flask contains 0.75 mol of carbon dioxide gas, CO2. What mass of carbon dioxide gas is in the sample ?
Solution
Practice Problems (text p. 186 text)
The triangle shown below is useful for problems involving number of moles, number of particles, and molar mass.
Converting from Mass to Moles
Write the formula used to convert mass to moles
Sample Problem
How many moles of acetic acid, C2H3COOH, are in a 23.6 g sample ? Solution
Practice Problems p. 187 text # 31 – 34
Sample Problem - Particles to Mass
What is the mass of 5.67 x 1024 molecules of cobalt (II) chloride, CoCL2
Solution
Practice Problems (text p. 190 # 35 – 38)
Sample Problem - Mass to Particles
Chlorine gas, Cl2, can react with iodine, I2 to form iodine chloride, ICl. How many molecules of iodine chloride are contained in a 2.74 x 10 –1g sample ? Solution
Practice Problems p. 191 text # 39 - 42
Ch. 5 Review p. 193 text
Mole Calculations Worksheet
1. Calculate the number of moles of the stated as substance of each of the following a samples
a 225 g of aspirin, C 9H8O4
b. 35 g of baking soda NaHCO3
c. 1.45 kG of gold
d. 0.84 g of hydrogen gas
e. 163 g sodium fluoride
f. 46.7 g of methanol, CH3OH
2. Calculate the mass of the following:
a. 12.4 mol of helium gas, He
b. 0.26 mol of butane. C4H10
c. 255 mol of ammonia
d. 1.8 mol of magnesium hydroxide
Answers
1a. 1.25 mol b. 0.417 mol c. 7.36 mool d. 0.42 mol
e. 3.88 mol f. 1.46 mol
2a. 49.6 g b. 15 g c. 4340 g d. 100 g
PECENTAGE COMPOSITION
(text p. 199 – 206)
State the law of definite proportions
carbon monoxide, CO, carbon dioxide, CO2,
C - 42.88 % by mass C – 27.29% by mass The percentage composition of a compound refers to the relative mass of each element in the compound.
Sample Problem
A compound with a mass of 48.72 g is found to contain 32.69 g of zinc and 16.03 g of sulfur. What is the percentage composition of the compound ?
• Mass percent of zinc =
• Mass percent of S =
Practice Problems . (text p. 201)
Calculating Percentage Composition from a Chemical Formula (p. 202 – 204) Percentage composition can be calculated from a known chemical formula. If you assume you have one mole of compound, you can use the molar mass of the compound, with its chemical formula, to calculate its percentage composition.
Example: Find the percentage composition by mass of HgS
Solution:
Sample Problem: Cinnamaldehyde, C9H8O, is responsible for the characteristic odour of cinnamon. Determine the percnetage composition of cinnamaldehyde by calculating the precents of carbon, hydrogen, and oxygen.
Solution:
Practice Problems (text p. 204 # 5 – 8)
Section Review p. 205 & 206
PERCENT COMPOSITION
WORKSHEET
Part A CALCUATING PERCENTS
When elements combine to form a given compound they are always present with same composition by mass
KClO3
% K % Cl % O
Fe(OH) 3
% Fe % O % H
Part B Calculating Simplest Formula (Empirical formula)
When elements combine to form compounds, they do so simple whole numbers.
Sample Problem
11.5 grams of sodium burdens in air to form 15.5 g of sodium oxide. Calculate the simplest formula of sodium oxide.
Na O mass # moles
Simplest mole ratio simplest formula Sample Problem
A compound consists of 27.29% carbon and 72.71% oxygen by mass. Calculate the simplest formula of this compound
C O
%
Mass in 100 g
# moles
Simplest mole
Ratio
Simplest formula Sample Problem
A compound consists of 52.17% carbon, 13.04% hydrogen and 34.78% oxygen. Calculate the simplest formula.
What does the empirical formula show?
What is the molecular formula of a compound?
Complete the following table:
Name of Compound
Molecular Formula
Empirical Formula
Lowest Ratio of Elements
Hydrogen peroxide
Glucose
Benzene
Acetylene (ethyne)
Aniline
Water
As you can see, it is possible for different compound to have the same empirical formula.
Determining a Compounds Empirical Formula (p. 208)
Sample Problem: Calculate the empirical formula of a compound that is 85.6 % carbon and 14.4% hydrogen.
Solution
Practice Problems p. 209 # 9 -15
Investigation 4-A (text p.p.212. 213)
Practice Problems (text p. 218)
7.1
STOICHIOMETRY
(text p.p. 234 – 259)
Balanced chemical equations are essential for making calculations related to chemical reactions.
Earlier, we learned how to classify types of chemical reactions and balanced chemical equations that describe them. We have also learned how to relate the number of particles and a substance to the amount of a substance in moles and grams.
Express Lab (text p. 235)
Complete the express Lab, Mole Relationships in a Chemical Reaction,
Answer Analysis 1 – 6
Particle Relationships in a Balanced Chemical Equation (text p.p. 235 – 237)
Recall that the coefficients in front of compounds and elements chemical equation tell us how many atoms and molecules participate in any reaction. For example, in the Haber
Process, where ammonia is prepared industrially from its elements:
N2(g) + 3H2(g) → 2NH3 (g)
This the equation tells us that one molecule of nitrogen gas reacts with three molecules of hydrogen gas to form two molecules of ammonia gas. You can use a ratio to express the number of atoms in the equation:
1 molecule N2 (g): 3 molecules H2(g) 2 molecules NH3 (g)
Example` If you wanted to 20 molecules of ammonia, how many molecules of nitrogen would you need?
Practice problems [text p. 237]
Mole Relationships in Chemical Equations
Until now, we have assumed that the coefficients in a chemical equation represent particles.
They can, however, represent moles. For example, from the Haber process shown earlier, shows that 1 mole of nitrogen reacts with 3 moles of hydrogen to form 2 moles of ammonia.
You can manipulate the mole ratios in the same way you manipulate ratios involving molecules. Sample Problem Using the equation above, how many moles are produced by 2.8 moles of hydrogen?
Practice problems [text p. 238)
Different Ratios of Reactants (text p. 239)
The relative amounts of reactant is are important. For example, carbon can combine with oxygen in 2 different ratios forming carbon dioxide, CO2, a product of cellular respiration in animals and one of the products of the complete combustion of a hydrocarbon fuel or
CO, a colourless, tasteless, odourless, highly poisonous gas that is responsible for hundreds of deaths in Canada and the United States each year.
Example Vanadium to form several different compounds with oxygen including V2O5, VO2, and V2O3. Determining the number of moles of oxygen that are needed to react with 0.56 mol of vanadium to form divanadium pentoxide.
Practice Problems [text p. 240]
Mass Relationships in Chemical Equations (text p. 241 – 246)
You have just learned that the coefficients in balanced equations represent moles as well as particles. Therefore, you can use the molar masses of reactants and products to determine the mass ratios for a reaction. Stoichiometric Mass Calculations
If you know what the amount of one substance in a chemical reaction [in particles, mass, or moles], you can calculate the amount of any other substance in the reaction to [in particles, mass, or moles], using the information in the balanced chemical equation.
Balanced chemical equations allow chemists to predict the amount of products that will result in a reaction involving a known amount of reactants. As well, chemists can use a balanced equation to calculate the amount of reactants they will need to produce the desired amount of products. They can also use it to predict the amount of one reactant they will need to completely react with another reactant.
Sample Problem: MASS TO MASS CALCULATIONS FOR PRODUCTS AND REACTANTS
Problem: Hydrazine, N2H4, and dinitrogen tetroxide, N204, formed a fuel mixture used to launch a lunar module. These two compounds react to form nitrogen gas and water vapour. If 50.0 g of hydrazine reacts with sufficient dinitrogen tetroxide, what mass of hydrogen gas is formed ?
Solution
i. Write a balanced chemical equation
ii. Convert mass of hydrazine to number of moles of hydrazine
iii. Calculate the number of moles of nitrogen, using the mole ratio of hydrazine to nitrogen iv. Convert the number of moles of nitrogen to grams
Practice Problems (text p. 246)
A GENERAL PROCESS FOR SOLVING STOICHIOMETRIC PROBLEMS
Step 1 Write a balanced chemical equation
Step 2 If you are given the mass or the number of particles, convert it to the number of moles
Step 3 Calculate the number of moles of the required substance based on the number of moles of the given substance, using the appropriate mole ratio.
Step 4 Convert the number of moles of the required substance to mass or number of particles as directed by the question
Sample Problem MASS AND PARTICLE STOICHIOMETRY
Problem: Passing chlorine gas through molten sulfur produces liquid sulfur dichloride. How many molecules of chlorine react to produce 50.0 g of sulfur dichloride ?
Solution
i. Write a balanced chemical equation
ii. Convert the given mass of sulfur dichloride to the number of moles
iii. Calculate the number of moles of chlorine gas using the mole ratio of chlorine to sulfur dichloride
iv. Convert the number of moles of chlorine gas to the number of particles of chlorine gas.
Practice Problems (text p. 248 bottom & 249 top)
Stoichiometry Problems Types 1-5
N2(g) + 3H2(g) → 2NH3 (g)
1. Mole -- Mole
Calculate the number of moles of ammonia [NH3 (g)] that can be produced from 5.6 moles of hydrogen gas [H2(g) ].
2. Mole – Mass
Calculate the number of grams of nitrogen gas (N2(g)) needed to react with five moles of hydrogen gas (H2(g))..
3. Mole – Volume (STP)
Calculate the number of moles of ammonia that would be produced from 8 litres of hydrogen gas STP
4. Mass – Mass Problems
a. Determine the mass of copper that can be produced from 217 g of copper (II) chloride
CuCl2 → Cu + Cl2
b. Iron reacts with oxygen gas to form iron (III) oxide. Determine the mass of iron (III) oxide produced from 112 grams of iron, given:
4 Fe + 3 O2 → 2Fe2O3
c. Calculate the mass of oxygen gas produced when 43.8 g of potassium chlorate is decomposed by heating, given:
2 KClO3 → 2 KCl + 3 O2
d. Calculate the mass of the BaSO4 produced from 24.1 grams of BaCl2 given:
BaCl2 + Na2SO4 → BaSO4 of + 2 NaCl
5. Mass – Volume Problems
a. Determine not the volume all of nitrogen gas at STP required to produce 12.6 g of ammonia, given N2 + 3H2 → 2 NH3
b. Determine the mass of aluminum required to produce 1.8 Litres of hydrogen gas at STP, given:
2 Al + 6HCl (aq) → 2AlCL3 + 3 H2
c. Magnesium reacts with HCl to produce hydrogen gas and magnesium chloride. Calculate the mass of magnesium required to produce 15 litres of hydrogen gas at STP , given:
Mg + 2 HCl → MgCl2 + H2
STOICHIOMETRY : TYPES 1-5
1. Fe + 2 HCl -------> FeCl2 + H2
(a) How many moles of iron would react with 6 moles of HCl ?
(b) How many moles of HCl would react with 0.1 moles of Fe ?
(c) How many grams of iron would produce 2 moles of H2?
(d) How many grams of iron would produce 300 grams of H2?
(e) How many grams of iron would produce 700 g of FeCl2?
(f) How many grams of iron would produce 22.4 L of H2 at STP ?
(g) How many grams of iron would produce 300 L of H2 at STP ?
2. C3H8 + 5 O2 ---------> 3 CO2 + 4 H2O
(a) How many moles of oxygen are required if 6 moles of C3H8 is to totally react?
(b) How many moles of CO2 will be produced is 15 moles of oxygen totally reacts?
(c) How many moles of water are formed if 12.2 moles of oxygen totally react ?
(d) What mass of oxygen will react with 2.5 moles of C3H8 ?
(e) What mass of carbon dioxide is formed if 30 grams of propane is totally reacted?
(f) What mass of carbon dioxide is formed if 263 grams of water is formed ?
(g) What mass of water is formed if 7800 grams of propane totally reacts ?
3. 3 H2SO4 + 2 Al ----------> Al2(SO4)3 + 3 H2
(a) How many moles of hydrogen gas are produced if 36.5 moles of sulphuric acid totally reacts?
(b) What mass of aluminum metal would produce 500 Litres of hydrogen gas at STP ?
(c) What mass of H2SO4 would produce 500 L of H2 at STP ?
(d) What mass of H2SO4 would produce 500 g of Al2(SO4)3 ?
(e) What mass of Al2(SO4)3 would be produced if 5786 L of H2 were produced at STP ?
7.2
THE LIMITING REAGENT
(p.p. 251 – 259 text) What are stoichiometric amounts?
In practice, however, there are often reactants left. In nature, reactions almost never have reactants in stoichiometric amounts. For example when an animal carries out respiration,
C6H12O6 (s) + 6O2 (g) → 6 CO6 (g) + 6 H2O (l)
there is an unlimited supply of oxygen in the air. The amount of glucose, C6H12O6 , depends on how much food the animal has eaten.
Determining the Limiting Reactant
What is meant by the limiting reactant?
Sample Problem IDENTIFYING THE LIMITING REACTANT
Problem Lithium nitride reacts with water to form ammonia and lithium hydroxide, according to the following equation:
Li3N (s) + 3H20 (l) → NH3 (g) + 3 LiOH (aq)
If 4.87 g of lithium nitride reacts with 5.80 g of water, find the limiting reagent.
Solution
i. Convert given masses into moles
ii. Use mole ratios of reactants and products to determine how much ammonia is produced by each amount of reactant.
•
• Amount of ammonia produced based on amount of H20
Alternative Method for Determining Limiting Reactant
Practice Problems (text p. 254 )
Inv 7-A (text p. 255)
The Limiting Reactant in Stoichiometric Problems
Since chemical reactions usually occur with one or more of the reactants in excess, it is often necessary to determine the limiting reactant before you carry out stoichiometric calculations.
Sample Problem THE LIMITNG REACTANT IN A STOICHIOMETRIC PROBLEM Problem White phosphorous consists of a molecule made up of four phosphorus atoms. It burns in pure oxygen to produce tetraphosphorus decaoxide.
P4 (s) + 5O2 (g) → P4O10 (s)
A 1.00 g piece of phosphorus is burned in a flask filled with 2.60 x 10 23 molecules of oxygen gas. What mass of tetraphosphorus decaoxide is produced?
Solution
i. Convert each reactant to moles and find the limiting reactant.
ii. Determine the limiting reactant
• Determine the amount of P4O10 that would be produced by 8.07 x 10 - 3 mol P4
• Determine the amount of P4O10 that would be produced by 0.432 mol O2
iv Using the mole to mole ratio of the limiting reactant to the product, determine the number of moles of phosphorus decaoxide that is expected
v. Convert this number of moles to grams
Or alternative solution
Practice Problems (text p.p. 257 & 258)
Section Review (text p.p. 258 & 259)
7.3
PERCENTAGE YIELD
(p.p. 260 – 270 text)
Theoretical Yield and Actual Yield
What does percent yield mean?
•
•
Write the formula used for calculating Percentage Yield
Sample Problem CALCUALTING PERCENTAGE YIELD
Problem: Ammonia can be produced by reacting nitrogen gas, taken from the atmosphere, with hydrogen gas.
N2 (g) + 3H2 (g) → 2NH3 (g)
When 7/5 a 10 1g of nitrogen reacts with sufficient hydrogen, the theoretical yield of ammonia is 9.10 g. If 1.72 g of ammonia is obtained by experiment, what is the percentage yield of the reaction ?
Solution
i. Divide the actual yield by the theoretical yield and multiply by 100 %.
Practice Problems (text p. 262)
Sample Problem PREDICTING ACTUAL YIELD BASED ON PERCENTAGE YIELD
Problem Calcium carbonate can be thermally decomposed to calcium oxide and carbon dioxide.
CaCO3(s) → CaO (s) + CO2 (g)
Under certain conditions, this reaction proceeds with a 92.4 % yield of ca;cium oxide. How many grams of calcium oxide can the chemist expect to obtain if 12.4 g of calcium carbonate is heated ?
Solution
i. Convert from grams of calcium carbonate to moles of calcium carbonate
ii. Use stoichiometry to calculate the theoretical yield of calcium oxide
iii. Convert from mol CaO to grams of CaO
iv. Calculate actual yield by multiplying theoretical yield by percentage yield
Or alternative method
‘
Practice Problems (text p. 264)
Inv 7-B (text p.266)
Complete Investigation 7-B, (text p. p. 255 – 267) and answer Analysis 1, 2, and
Conclusion 3
STOICHOMETRY - TYPES 6-8
6. VOLUME-VOLUME (at the same conditions)
What volume of N2 at 20oC and 730 mm Hg would produce 400 Litres of NH3 at the same conditions?
N2 + 3 H2 ------> 2 NH3
7. VOLUME - MASS (not STP)
What mass of Al is required to form 1800 Litres of H2 at 30oC and 80 kPa ?
2 Al + 3 H2SO4 ------> Al2(SO4)3 + 3 H2
8. MASS- VOLUME (not STP)
What volume of hydrogen gas is formed at 80oC and 0.5 atm if 130 grams of Al metal is reacted?
2 Al + 3 H2SO4 ------> Al2(SO4)3 + 3 H2
SCH3U STOICHIOMETRY: TYPES 6-8
1. 4 Li + O2 ---------> 2 Li2O
What mass of Li would be needed to react with 35 L of oxygen gas at 25 degrees C and 900 mm Hg ?
2. Solid ammonium dichromate, when heated, yields gaseous nitrogen, chromium (III) oxide and water.
(NH4)2Cr2O7 ---------> 4 N2 + Cr2O3 + 4 H2O
What volume of nitrogen gas at 80 degrees C and 104.6 kPa would be formed if 8 grams of water is also formed ?
3. 2 BBr3 + 3 H2 --------> 2 B + 6 HBr
(a) What volume of boron bromide at 300 degrees C and 65 kPa would produce 600 L of hydrogen bromide at the same conditions ?
(b) What mass of boron is formed if 3700 L of hydrogen gas at 1500 degrees C and 1.8 atm is totally reacted ?
4. 2 C10H20O + 29 O2 --------> 20 CO2 + 20 H2O
(a) What volume of oxygen gas is required at 300 degrees C and 800 mm Hg to produce 2500 grams of carbon dioxide ?
(b) What volume of water vapour is produced at 300 degrees C and 85 kPa if 1800 grams of oxygen gas totally reacts ?
(c) What volume of carbon dioxide at 400 degrees K and 1.75 atm will be produced if 2500 grams of C10H20O totally reacts?
Name __________________________
A Selection of Enjoyable Stoichiometry Problems
Consider the combustion of ethyl alcohol (C2H5OH) being investigated by a scientist called Barbara.
C2H5OH (l) + 3O2 (g) → 2CO2 (g) + 3H2O(g)
1. How many moles of carbon dioxide are produced when Barbara burns 1.50 moles of C2H5OH and
2. What mass of carbon dioxide is produced 1.5 moles ethanol is burned?
3. What mass of carbon dioxide is produced 23.0 in g of ethanol is burned?
Consider the following reaction for questions 4, 5 and 6
3Ag(s) + 4HNO3(aq) → 3AgNO3 (aq) + NO(g) + 2H2O(l)
4. How many moles of NO are produced when 1.5 moles of Ag reacts with excess HNO3?
5. How many grams of NO are produced when 1.5 moles of Ag reacts with excess HNO3(aq)?
6. How many grams of AgNO3 are produced when 189 g of HNO3 reacts with excess Ag?
Consider the following reaction:
3Cu(s) + 8HNO3 (aq) → 3Cu(NO3) 2(aq) + 2NO(g) + 4H2O(l)
7. How many moles NO are produced the reaction 6.35 g of Cu with excess nitric acid?
8. What mass of water is produced 12.6 g of nitric acid reacts with excess Cu?
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Thousands of chemicals have been introduced into our society over the last one hundred years. From furniture to the makeup and shaving cream used every day, nearly everything has at least one chemical proven to cause health issues. These new chemicals, which were not around hundreds of years ago, are exposed to our body’s daily. Forty-two billion pounds of these dangerous chemicals enter American commerce daily. The film moves from presenting health issues to showing how there is a movement towards solving some of them. Green chemistry has grown greatly and is a step towards solving this issue. Green chemistry was a 2.7 billion-dollar industry in 2015. Many companies, of all kind, use green chemistry to reduce harmful chemicals in their products. Although some companies are moving towards a possible solution to this question, many companies are not and the fight for healthy products continues. In conclusion, the documentarian’s use of rhetorical devices was effective and persuades the viewer to recognize the issue regarding the millions of dangerous chemicals in products we use daily. If companies were to use healthier alternatives in their products, in addition to tougher legislation limiting which chemicals are allowed on the market, the issue could be…
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