Chemistry Gr.11

Chemistry, 2014 UNIT 1 – MATTER TRENDS AND CHEMICAL BONDING

History of the atom
DEMOCRITUS
Smallest particle of matter is called an atom
Atoms are in constant motion and have empty space between them
ARISTOTLE

4 element theory of matter (earth, air, water, fire)
Had different combinations of the 4 qualities: hot, cold, dry, moist
JOHN DALTON

Expanded on Democritus’s theory:
All matter is composed of tiny, indivisible particles
All atoms of an element have identical properties
In chemical reactions atoms join together/separate but are not destroyed
Atoms of 2 or more elements can combine in ratios to form new substances
Law of constant composition
J.J THOMSON

Hypothesized existence of electrons
Negatively charged electrons are distributed inside the atoms which is a positively charged sphere consisting of mostly empty space (Plum Pudding model)
ERNEST RUTHERFORD
Gold foil experiment: shot alpha particles through very thin god foil
Most would travel through but some deflected
Atoms contain positively charged core surrounded by empty space containing electrons
Only a small portion volume could be attributed to nucleus
JAMES CHADWICK
Nuclei must contain neutral and positive particles
An atom is composed of a nucleus complaining protons and neutrons

Writing and balancing chemical equations
1. Check for diatomic gases and place the subscript 2 (N, O, F, Cl, Br, I, H)
2. Balance formulas of compounds and place c
3. Conservation of atoms (both sides should have equal # of ions of each element)
4. Law of conservation of mass (both sides should have equal mass)

Isotopes
2 or more forms of the same element that differ in their mass number because they have different numbers of neutrons in their nuclei (atomic # will still be the same)

Ex. U-235 235 - 92 = 143NO
Mass: 235 Protons – Mass = Neutrons
92P+

Isotopes of Hydrogen
Hydrogen –1 proton, no neutrons
Deuterium (heavy water) – 1 proton, 1 neutron
Tritium – 1 proton, 2 neutrons
Average Atomic Mass
Units = (unified atomic mass)

Ex. Boron consists of 2 isotopes: B-10 at 19.78% B-11 at 80.22% = 10.8µ
Check isotope problem set for practice

Mass Spectrometer
Stage 1: Ionization
The atom is ionized by knowing out 1 or more electrons to give a positive ion. Mass spectrometers always work with positive ions. ONLY electrically charged particles are affected by the magnetic field.

Stage 2: Acceleration
Ions are accelerated so they all have the same kinetic energy.

Stage 3: Deflection
The ions are deflected based on 2 things
a) Mass – the lighter they are the more they are deflected
b) Charge – ions with a higher charge are deflected more

Stage 4: Detection
The beam of ions is detected on a photographic plate or by a computer and a film or printout is made

Mass spectrometer film/diagrams
Start with the bigger mass with a +1 charge, then smaller mass with +1, and then bigger mass with +2, and the smaller mass with +2
Ex. K-41 K-39 K-41 K-39 +1 +1 +2 +2
The ions with a larger percent abundance should be thicker lines, and the smaller percentages should have thin lines

Radioisotopes
An unstable isotope that emits radiation while decaying to become stable

Alpha particle
Beta particle
Gamma particle
Another name
Helium nucleus
High energy electron
High energy electromagnetic waves
The symbol for this particle α β γ How nucleus of a radioisotope is altered by emission of this particle
Decreases the mass # by 4 and the atomic # by 2
+2 charge
Converts a neutron into a proton
Increases atomic # by 1
-1 charge
No change in mass number or atomic number

Penetrating ability
Penetration in air is a few centimeters. Skin can stop it.
In air is a few meters.1-2mm piece of paper can stop it
In air is unlimited. 1m of lead or concrete can stop it

Alpha Decay Equation
The reason alpha decay occurs is because the nucleus has too many protons which cause excessive repulsion. In an attempt to reduce the repulsion, a Helium nucleus is emitted (2 protons, 2 neutrons)
Beta Decay
Beta decay occurs when the neutron to proton ratio is too great in the nucleus and causes instability. In basic beta decay, a neutron is turned into a proton and an electron.
Gamma Decay
Gamma decay occurs because the nucleus is at too high an energy. The nucleus falls down to a lower energy state and, in the process, emits a high energy photon known as a gamma particle.

Half Life
Time required for half of a sample to decay
Half-life occurs naturally in some of the radioactive elements while it could be artificially stimulated in some other elements

Carbon Dating (creds to Ravneet)
At this moment, your body has certain percentage of carbon – 14 atoms in it, and all living plants have the same percentage. As soon as organisms die, it stops taking in new carbon. The carbon –14 decays with its half-life of 5700 years, while the amount of carbon –12 remains constant in the sample.
By looking at the ratio of carbon – 12 to carbon – 14 in the sample and comparing it to the ratio in a living organism, it is possible to determine the age of a formerly living thing

Transmutation
Reaction in which new identity of nucleus is formed (i.e. new element)
Often a nucleus is struck by a neutron or another nucleus e.g.
Ex.

Nuclear Fission
When lighter atoms combine to form heavier ones
Need extremely high temperatures and speed
Only occurs naturally in the sun and stars
For practice with fore mentioned, see page 219 #3-6 and 223 #3-6 in your textbook

CANDU – Page 34-35 #20-23 (probably gonna be a multiple choice question on this idk)

Energy levels
When energy (heat/electricity) is applied to atoms, excited electrons gain a certain quantity of energy. With this energy they jump from a lower to a higher energy level. As they drop back to lower energy levels, they release energy in photons corresponding to a few precise wavelengths.

Spectra
Violet light: 62
Indigo light: 52
Green light: 42
Red light: 32

Quantum Theory
Orbitals and sublevels
Number of sublevels an energy level can contain = number of sublevels same as n
Second energy level – s, p (2 sublevels)
Third energy level - s, p, d (3 sublevels)
Fourth energy level - s, p, d, f (4 sublevels)

Bohr’s Atomic Structure Configuration:
1st energy shell: 2 electrons 2(1) 2 = 2
2nd energy shell: 8 electrons 2(2) 2 = 8
3rd energy shell: 18 electrons 2(3) 2 = 18
4th energy shell: 32 electrons 2(4) 2 = 32

Heisenberg’s uncertainty principal: states that It is impossible to know precisely both the speed and location of an electron at the same time. Bohr’s model said we could (it was wrong.)

Electron configuration (atomic #)

1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10
Short form: Go to last noble gas and list electrons from there
Ex. Ca = [Ar] 4s2
When removing electrons, always remove from last ENERGY LEVEL
For practice on electron configuration see here
Isoelectronic Ions
Atoms with the same number of electrons

Electron Promotion
Exceptions to explain some chemical results and increase stability
Stabilizes the d orbital (should be full or half full)
Ex. Silver – Ag
[Kr] 5s2 4d9  it’s worth it to promote an electron to fill the d orbital
[Kr] 5s1 4d10  5s promoted the electron to 4d to stabilize

Radon
Radon is a radioactive gas found in the environment. It is produced by the decay of uranium. It is odorless, tasteless and emits radiation.
How Radon Enters a House
Radon escapes from bedrock into outdoor air. Radon can escape into a building if it is built ontop of bedrock or soil that contains uranium. Mostly, it enters through cracks in the foundation, walls and floors of basements.
Health Effects of Radon
It is the second leading cause of lung cancer after smoking. After breathing in radon gas, they emit “alpha particles” resulting in lung cell damage. The main source of Radon entrance in a house is through soil.
Removing Radon
Common detectors: Alpha Track Detection & Electret Ion Chamber. The most common way is Active Soil Depressurization: a hole is drilled in basement & a pipe draws radon gas from under the house & pushes it outside.
Radioactivity
When solid Radon decays to form gas, it loses 2 protons and 2 neutrons. These particles are known as alpha particles. This element is radioactive because it decays, losing an alpha particle and forming polonium. Radon decays to form radioactive alpha particle:

Periodic Trends
Atomic Radius
The distance from the center of the nuclei to the outermost valence electron, controlled by:
1. Number of shells
2. Nuclear pull (strength of attraction felt by valence electrons)
3. Shielding effect (electrons in innermost shells block outermost electrons from nucleus’ attractive force)
Top to bottom:
More reactive
Softer
More metallic
Ionization energy decreases
Electron affinity decreases
Electronegativity decreases
Nuclear pull decreases
Atomic radius increases

Nuclear charge
Measure of the strength of a nucleus’s pull (the more protons, the stronger the nucleus)
Positive ions are smaller than neutral atoms  electrons are pulled in
Negative ions are larger than neutral atoms  electrons are repelling

Ionization energy
The amount of energy required to remove an electron from at atom/ion in the gaseous state. Depends on which electron is being removed (more difficult to remove electrons closer to nucleus) Most weakly held electron will be removed first (valence shell electrons) Write out electron configuration to help explain
1st I.E.  energy required to remove 1st valence electron, and so on
Electron Affinity
Energy change that accompanies the addition of an electron to a gaseous atom (The more it wants an electron smaller it’s size)
Most electron affinities are negative because most elements release energy when they become negative ions
E.g. F(g) + e-1  F-1 + Energy

Lewis Diagrams for ionic bonding
The Lewis structure for the ionic compound LiF would be: + 

Lewis Structures (Covalent)
1. Calculate the number of total electrons
2. Arrange atoms around central atom (highest bonding capacity)
3. Distribute them to satisfy the octet rule
4. Draw lines between the elements that share electrons (ex. NF3)

Bonding Capacity
For groups 1, 2 and 3, BC = group #
For groups 4 to 8, C = 8 - # of valence electrons

Multiple Bonding
Sometimes the only way that bonding atoms can attain a stable octet is by sharing more than one pair of electrons
The ONLY elements that can form multiple covalent bonds with itself or each other are: C, N, O, S, P

Molecules that do not obey the octet rule
BF3 (boron will have less than 8 electrons in its valence shell)
SF4 (sulfur will have more than 8 electrons in its valence shell)
Why can sulfur EXPAND its octet?
We need to look at the electron orbital configuration for sulfur
Sulfur [Ne]3s23p4
It only needs 2 more electrons to fill its valence shell
BUT, when n=3, this energy level contains 3s, 3p AND 3d sublevels (18 electrons total)
Sulfur can use its EMPTY 3d orbitals to hold EXTRA electrons
A similar explanation can be used for Phosphorus (e.g. PCl5)

Coordinate covalent bonding
A covalent bond in which both shared electrons comes from one atom
The other atom does not provide any electrons (ex. NH4)

Electronegativity
An actual number that describes the ability of an atom to attract electrons, measured on a scale of 0.0 to 4.0
When electrons are shared by two atoms a covalent bond is formed.
When the atoms are the same they pull on the electrons equally. Example, H-H.
When the atoms are different, the atoms pull on the electrons unevenly. Example, H-Cl

Polar bonds
A bond that has an uneven distribution of charge due to an unequal sharing of bonding electrons Ex. HCl
Polar bonds contain dipoles determined by electronegativity (positive and negative ends)
Higher electronegativity is the negative side
Subtract the two electronegativity’s to determine the type of bonding
For a molecule to be polar it must meet 2 requirements
Contains at least 1 polar bond
Asymmetrical

It is possible for a molecule to contain polar bonds but be a non-polar molecule if it is symmetrical (CCl4)

Intramolecular forces
Bonds holding atoms together to form a molecule (H2O or NaCl)
Ionic – a metal and a nonmetal
Covalent – 2 non-metals
Intermolecular forces
Forces between molecules, strong forces cause high melting point, boiling point, freezing point, and hardness
Hydrogen bonding (strongest)
Special dipole-dipole bond between H and another extremely electronegative atoms Flourine, Oxygen, Nitrogen
Dipole-dipole
The negative end of one polar bond is attracting the positive end of another polar bond (Ex. HCl)
London dispersion (weakest)
Very weak force between non-polar molecules
For a moment, electrons are located on one side of the atom resulting in an “instantaneous dipole”
The negative end attracts the positive end of another atom

UNIT 2A – NOMENCLATURE
Naming binary ionic compounds
Metal + Non-metal, change non-metal suffix to –ide (ex. NaCl  Sodium ChlorIDE)
Criss cross rule – exchange subscripts

Binary molecular compounds
Non-metal + non-metal
Add correct prefix before each element

Radicals
A group of covalently bonded atoms that has a net charge (positive or negative)

Common Radicals
NO31-
Nitrate
ClO31-
Chlorate
CO32-
Carbonate
SO42-
Sulfate
PO42-
Phosphate

Hydrogen Radicals
Adding H+ to radicals causes the charge on the radical to change
HCO31-
Hydrogen carbonate or bicarbonate
HSO41-
Hydrogen sulfate or bisulfate
HSO31-
Hydrogen sulfite or bisulfite
HPO42-
Monohydrogen phosphate
H2PO41-
Dihydrogen phosphate

Examples
NaC2H2O3
Sodium acetate
KOH
Potassium hydroxide
Sr(HCO3)2
Strontium bicarbonate
Ca(MnO4)2
Calcium permanganate
e.g. K+1 and HCO31-
KHCO3  potassium bicarbonate or potassium hydrogen carbonate
Common Radical prefix/suffixes
Per____ate
_____ates
______ites
hypo_____ite
NO41-
ClO41-
CO42-
SO52-
PO53-

NO31-
ClO31-
CO32-
SO42-
PO43-
NO21-
ClO21-
CO22-
SO32-
PO33-
NO1-
ClO1-
CO2-
SO22-
PO23-

 Add 1 Oxygen  Remove 1 Oxygen  Remove 2 Oxygens

Examples
KNO3
Potassium nitrate
KNO2
Potassium nitrite
Mg(NO)2
Magnesium hyponitrite
Mg(NO4)
Magnesium pernitrate
e.g. Li+1 and SO5-2
Li2SO5 Lithium persulfate

Multivalent compounds
Compounds with metals that have more than one charge (Ex. FeSO4 vs. Fe2(SO4)3)

Ion Stock System Traditional System Fe 3+ iron (III) ferric Roman Numerals Fe 2+ iron (II) ferrous Cu 2+ copper (II) cupric Cu 1+ copper (I) cuprous Co 3+ cobalt (III) cobaltic Co 2+ cobalt (II) cobaltous Sn 4+ tin (IV) stannic Sn 2+ tin (II) stannous Multivalent Ionic Compounds with Radicals Pb 4+ lead (IV) plumbic Pb 2+ lead (II) plumbous Hg 2+ mercury (II) mercuric Hg2 2+ mercury (I) mercurous
Naming Acids
1. Binary Acids
Only H and another element are present
Ex: HBr (aq) (means HBr dissolved in water)
Prefix: hydro- Suffix: -ic
Hydrobromic acid
H2S (aq) = Hydrosulfuric acid

2. Oxyacids
Anion (metal) contains oxygen
“ous” and “ic” endings apply here
**do not use hydro
Ex. Ex: H2SO4(aq) sulphuric acid H2SO3 (aq) sulphurous acid

Naming Hydrates
Compounds with water in their structure
This is different from aqueous because water is part of the molecule
H20 can usually be removed when heated
Ionic compounds that absorb water have a water molecule attached to them and are called HYGROSCOPIC (meaning they want to absorb water)
ANHYDROUS: without water
HYDRATED: with water
Prefixes describe how many H2O molecules are attached
Ex. CuSO4•5H2O is copper(II) sulfate pentahydrate.

UNIT 2B – CHEMICAL REACTIONS
Types of reactions
1. Synthesis
Simple substances combine to form more complex compounds
Product can either be ionic or molecular A + B  AB (more complex) Ex. 2Na + Cl2  2NaCl
Element + element  compound
Compound + element  complex compound
Compound + compound  very complex compound

2. Decomposition
A compound is broken down into simpler substances
AB  A + B
Ex. 2HCl  H2(g) + Cl2 (g)
Special case: Metal CARBONATES decompose to produce the metal oxide and carbon dioxide gas.
Example: MgCO3 (aq)  MgO(s) + CO2 (g)

3. Single displacement
An element reacts with a compound and displaces a second element from the compound to form an element and a new compound
AB + C  AC + B
*Don’t forget to follow the activity series. Some reactions will not occur.
C has to be more reactive than A for a reaction to occur

Single displacement usually occurs between…
a) an active metal + an acid
When a metal which is above hydrogen in the activity series is reacted with an acid, hydrogen is liberated and a salt is formed. Zn + 2HCl  H2 + ZnCl2 Mg + H2SO4  H2+ MgSO4
b) A metal + a salt
Each metal in the activity series displaces any metal below it to form a salt in solution.
Cu + AgNO3  CuNO3 + Ag
Fe+ CuSO4  FeSO4 + Cu
c) A Halogen + halide salt
A halogen will displace any less active halogen from a halide salt. The order of activity decreases going from top to bottom down the halogen family in the periodic table. Cl2 + 2NaI  2NaCl + I2

4. Double displacement
Precipitation reactions – formations of a solid
In order to determine which one of the products will precipitate, see solution table (REMEMBER SUBSCRIPTS)
Neutralization reactions – formation of water + a salt
Also called acid-base reactions, reaction between an acidic compounds and a basic compound pH = 7 if all acid/base reacts
Use bro. blue as indicator – should turn green
HCl(aq) + NaOH(aq)  H2O + NaCl(aq)
H2SO4(aq) + KOH(aq)  2H2O + K2SO4(aq) AB + CD  AD + CB
The valence of multivalent elements stay the same on both sides of the equation*

5. Combustion – burning of a hydrocarbon
A hydrocarbon is a molecular compound containing only carbon, hydrogen and sometimes oxygen.
Complete combustion of a hydrocarbon always produces only 2 products: CO2 and H2O
Example: Write a balanced chemical equation for the combustion of propane (C3H8). C3H8 + 5O2  3CO2 + 4H2O

UNIT 3 – QUANTITIES IN CHEMICAL REACTIONS
Scientific Notation
Used when working with very large/small numbers
Ex. 0.0052 move decimal 3 places to right  5.2 x 10-3 0.00065  move decimal 3 places to right  6.5 x 10-4
Base number is always written as a single digit
Significant digits
Any digit of a number that is known with certainty
All of the digits from 1 to 9 are significant
1.234g has 4 sig figs 1.2 g has 2 sig figs
Zeros in between non-zero digits are significant
1002 kg has 4 sig figs 3.07 mL has 3 sig figs
Any zero to the right of a nonzero digit (trailing zeroes) is significant 0. 230 mL has 3 sig figs 0.40 has 2 sig figs
- All zeros to the left of the first non-zero digit (leading zeros) are NOT significant. They merely mark the position of the decimal point 0.001 degrees C has only 1 sig fig 0.012 g has 2 sig figs
Space holding zeros on numbers less than one (they are after the decimal but before non-zero digits are Not significant. 0.023 has 2 sig figs (write as 2.4 x 10-2)
Stoichiometry (this will basically dictate whether you pass or not)
The study of the quantitative aspects of chemical reactions
The Mole (mol)
One mole is defined as 6.02 X 1023. (3 sig figs) also known as Avagadro’s constant (NA)
Molar Mass (g/mol)
Equal to average atomic mass found on periodic table (round to 1 decimal place)
Na = 23.0 g/mol
Mg = 24.3 g/mol
O2 (there are TWO moles of O) = 32.0 g/mol
1 mol of CaCl2 = 111.1 g/mol 1 mol Ca x 40.1 g/mol
+ 2 mol Cl x 35.5 g/mol = 111.1 g/mol CaCl2

Mole formula #1

Mole Formula #2

Percent composition Formula

Empirical Formula (create chart)

Molecular Formula
To determine molecular formula you must know
Simple formula
Molar mass of compound

Percent Yield
The expected or calculated amount of product formed is called the theoretical yield.
The amount of product actually formed in the experiment is called the actual yield.

Percent yield tells us how “efficient” a reaction is. Percent yield cannot be bigger than 100%.
To find theoretical yield, calculate moles then convert to mass and divide by actual yield.

Limiting and Excess reagents
You will know it is a limiting question if you are given the mass of 2 reactants. The problem is that you must determine which of the reactants will be consumed first, limiting the amount of product made. It will be called the LIMITING REAGENT. The other reactant will not be totally consumed so it is in EXCESS.
1. Write your equation and balance
2. Find moles of both reactants
3. Divide moles by the coefficient for that substance to determine the number of cycles that can be completed
4. Whichever number is smaller is the LIMITING reagent
5. Use the limiting reagents MOLES to find the number of moles of product by multiplying it by the unknown/known ratio and convert to mass

UNIT 4 – SOLUTIONS AND SOLUBILITY
Components of a solution
Solute – the substance that is being dissolved in a solvent *e.g. sugar
Solvent – the medium in which a solute is dissolved (often a liquid) e.g. water
When water is used as a solvent, the solutions are called aqueous solutions
e.g. NaCl(aq) - solid in a liquid C2H5OH(aq) - liquid in a liquid

I2(al) – iodine dissolved in alcohol
When a solid dissolves in a liquid solvent we say that it is SOLUBLE in that solvent
If the solute does not substantially dissolve in a given liquid we say that it is INSOLUBLE
When two liquids dissolve in each other they are said to be MISCIBLE (usually the liquid present in the smallar quantity is regarded as the solute)
IMMISCIBLE liquids will not dissolve in each other
Electrolytes – solutions that conduct electricity because they contain ions
Non-electrolytes – do not conduct electricity
Acidic – pH < 7, contain H+ ions
Litmus paper  red
Phenolphthalein  stays clear
Bromthymol blue  yellow

Water as a universal solvent
Ionic compounds are very soluble in water
Ionic molecules dissociate or separate into aqueous ions
These ions are now SOLVATED
The process by which ions are attracted to and surrounded by solvent molecules
If the solvent is water the term used is “hydrated”

Dissolving non-polar molecules
“Like dissolves like”
Non polar molecules are held together by weak London forces, but one weakly held non-polar substance could go between another weakly help non-polar substance
Solubility is polarity dependent
Consider a molecule of ethanol  ∴ It can form Hydrogen bonds and dipole-dipole bonds with water, so it is SOLUBLE in water

Gasoline is OCTANE – CH3CH2CH2CH2CH2CH2CH2CH3
NON-polar molecule INSOLUBLE in water
It is not strong enough to break forces between water molecules
Note: some ionic compounds insoluble in water
Mostly carbonates, phosphates and sulphites
Why? They have higher charges (on ions) so they are more difficult to pull apart – stronger ionic bonds

Calculating concentration
Concentration – the quantity of a given solute in a solution
Dilute – a relatively small amount of solute per unit volume of solution
Concentrated – a large amount of solute per unit volume
Concentrations can be calculated and expressed in numerous ways: v/v, w/w, w/v, ppm v/v  volume by volume Percent concentration w/w  weight by weight w/v  weight by volume ppm  parts by million

PPM
For very dilute solutions, concentrations are sometimes expressed in parts per million, abbreviated as ppm
These questions can get very confusing. I have found that the best solution is to memorize the equivalent values for ppm and then just do unit conversions.

1 ppm is equal to : = 1g/m3 or 1g/1000L = 1mg/L = 1 mg/kg

Molar concentration (Mole formulas #3)

[ ] square bracket also a symbol of concentration
e.g. [HCl(aq)] means “the concentration of HCl……”

Chemical, Ionic, and Net Ionic equations
CHEMICAL EQUATIONS are written as if all substances were molecular, even though some substances may exist as ions.
HCl (aq) + NaOH (aq)  NaCl (aq) + H2O (l)

IONIC EQUATIONS have the substances which exist as ions (i.e. dissociate in water) written in ionic form (solids and liquids do not dissociate)
H+ (aq) + Cl- (aq) + Na+ (aq) + OH-(aq)  Na+ (aq) + Cl- (aq) + H2O (l)

NET IONIC EQUATIONS are ionic equations with the Spectator ions removed. H+ (aq) + OH-(aq)  H2O (l)

SPECTATOR IONS do not participate in a reaction (that is they do not react to form a new substance). Common Spectator ions are Group I, many Group II, and NO3- (nitrate) and C2H3O2- (acetate) ions.

Solubility curves
Unsaturated Solution- more solute can still be dissolved if the temperature remains constant
Saturated solution – the maximum amount of a solute that can be dissolved in the solvent has been added. No more solute can be dissolved if temp remains constant.
Supersaturated solution – More than the maximum amount of a solute that can be dissolved in the solvent has been added. It is an unstable situation

Solubility
The concentration of a saturated solution of a solute in a solvent at a specific temperature and pressure
Units are in g/100 mL
Temp and pressure must be stated because they both affect solubility
Solubility is a characteristic physical property of a substance

Reading a solubility chart
The curve shows the # of grams of solute in a saturated solution containing 100 mL or 100 g of water at a certain temperature.
Any amount of solute below the line indicates the solution is unsaturated
Any amount of solute above the line in which all of the solute has dissolved shows the solution is supersaturated.
If the amount of solute is above the line but has not all dissolved, the solution is saturated and the # grams of solute settled on the bottom of the container = total # g in solution – # g of a saturated solution at that temperature. (according to the curve)
A concentration directly on the line represents the saturation point
Solutes whose curves move upward w/ increased temperature are typically solids because the solubility of solids increases with increased temperature.
Solutes whose curves move downward with increased temperature are typically gases because the solubility of gases decreases with increased temperature
For more practice see this worksheet

Select Precipitate
Prepare a chart to help identify precipitates that will form
Ex. Ba(NO3)2, Mg(NO3)2, and AgNO3
Nitrate is a spectator ion so it can be ignored
Determine the order that we add these solutions, so that we have only one precipitate that occurs with the each addition.
"ppt" indicates an insoluble precipitate; "sol" indicates a soluble compound that remains dissolved as ions
Ba2+
Mg2+
Ag+
SO42- ppt sol ppt Cl- sol sol ppt OH- sol ppt ppt Acids and Bases (creds to Ravneet Gill✌) Properties
Acid
Base
Produce hydronium ions in water
Taste sour
React with metal to form hydrogen gas
Form electrolytes
React with base to form water and salt
Turn blue litmus paper red
Phenothalien clear
Bromthymol blue remains clear pH less than 7 produce hydroxide ions in water
Taste bitter
React with acids to form water and salt
Are electrolytes
Turn red litmus paper blue
Turn Phenothalien pink
Bromthymol blue remains blue pH greater than 7

Ionization of Acids
H+ or H3O+ (hydronium) is produced in water. Acids dissolve and ionize into hydrogen ions and negative ions in solution. Is the process by which a neutral atom/ molecule is converted into an ion
Polyprotic equations: when there is more than 1 H to lose or receive, only 1 can be removed at a time
Example:
H2SO4 + H2O  H+(aq) + HSO4 -
HSO4-+ H2O  H+(aq) + SO4 -2
Dissociation of Bases
The separation of ions that occurs when an ionic compound dissolves in water
They break apart into positive and negative ions in solution
Example:
Ba(OH)2  Ba2+ (aq) + 2OH- (aq)
Arrhenius Acids & Bases
Acid: in aqueous solutions form hydrogen or hydronium ions (H3O+)
Base: in aqueous solutions form hydroxide ions (OH-)
Bronsted - Lowry Acids & Bases
Bronsted – Lowry Acid: proton donor.
Bronsted – Lowry Base: proton acceptor.
This means that the proton moves from the acid to base
Cannot have 2 acids or 2 bases on one side

Amphriprotic or Amphoteric is a substance that can be used as an acid or base defending on what it is being mixed with (conditions)
Example: Water

pH Scale
Indicates the concentration of H+ ions present in a solution
More H3O+ ions = acidic, less=basic
Significant digits are AFTER decimal point
To convert concentration of Hydronium to pH use negative log
Examples E = 10 to the power of --
[H3O]+
pH
Typed into Calculator
6.25 x 10 -5
4.204
-log (6.25x10 -5)
To convert pH to concentration of Hydronium use 10 – pH
Unit for concentration = M (mol./grams)
Examples
[H3O]+ pH Typed into Calculator
6.25 x 10 -5
4.204
10 ^ -4.204

Titration & Neutralization Reactions
Example of Neutralization:
LiOH(aq) + HNO3(aq) è LiNO3(aq) + H2O(l)
Ionic equation: Li + + OH - + H+ + NO3 -è Li+ + NO3- + H2O(l)
Net ionic equation: OH - + H+ è H2O(l)
Titration: procedure involving the carefully measured and controlled addition of a solution from a buret into a measured volume of a sample solution
Neutralization: double displacement reaction between an acid and base producing water and a salt
Standard Solution: a solution of precisely and accurately known concentration
Endpoint: the point in a titration at which a sharp change in a property occurs (e.g., a colour change)
Titrant: the solution in the buret during a titration
Indicator: a compound that changes color at a specific pH value or in the presence of a particular substance
Strong Acid: an acid that ionizes almost completely (>99%) in water to form aqueous hydrogen ions
Strong Base: an ionic hydroxide that dissociates 100% in water to produce hydroxide ions
Burette (Buret): graduated glass tube with a tap at end for delivering known volumes of a liquid
Equivalence Point: point at which chemically equivalent quantities of an acid/base have been mixed, can be found by an indicator.
Stoich Problems
1. Write balanced chemical equation
2. Write information under each reactant. Convert mL to L
3. Determine the number of moles of given reactant.
4. Write a mole ratio & solve for moles of unknown acid/base
5. Use c = n/V to find concentration of unknown acid or base

Conjugate acid base pairs: base formed by removing the H+ ion from an acid
Example

HCO3 + H2O  H2CO3 + OH B A A B

You will only actually understand how to do stoich once you try some problems.
Do not rely on these notes to get you through.

UNIT 5 – GASES
KMT and States of matter

Forces between molecules
Entropy/degree of disorder
Types of motion
Solid
Ionic bonds
Covalent bonds
Metallic bonds
Some dipole-dipole
High mp/bp
Don’t flow or compress
Ordered
Low entropy
Vibrational
Liquid
Dipole-dipole
Hydrogen
Flows
Less orderly
Rotational
Vibrational
Translational
Gas
Noble gases
London forces only
Diatomic
Polar
No hydrogen bonds because distance is too far
Expand to take shape of container
Easily compressed
No organization
Greatest entropy
Rotational
Vibrational
Translational

Kinetic molecular theory: the idea that all substances contain particles that are in constant random movement and colliding with other particles

Pressure
*Standard pressure (S.T.P.) = 101.325 kPa (0oC) = 1 atm = 760 mmHg = 1 torr

*Standard ambient = 100 kPa Temperature & pressure (S.A.T.P.) (25OC)

Temperature conversions
Convert Celcius to Kelvins (OC + 273 = temperature in Kelvins)
Temperature should always be expressed in Kelvins for any calculations

Absolute zero
Charles found that the point where a gas would have no volume would be -273 degrees Celsius

Boyles Law

As the pressure on a gas increases, the volume of the gas decreases proportionally
As the pressure on a gas decreases, the volume of the gas increases proportionally
Temperature is constant

Charles Law

If the temperature increases, the volume increases; if the temperature decreases, then the volume decreases.
Pressure is constant

Gay-Lussac Law

Combined gas law

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Ideal gas Law R= ideal gas law constant (8.314)

Properties of ideal gases
A gas that obeys all the gas laws perfectly under all conditions
V-T and P-T graphs are perfectly straight lines
Gas does not condense to a liquid when cooled
Gas volume = 0 at absolute zero

Dalton’s law of partial pressure
The total pressure exerted by a gas mixture is the sum of the partial pressures of the gases in that mixture. PT = P1 + P2 + P3 + .....

Gay Lussac’s Law of combining volumes
When measured at the same temperature and pressure, volumes of gaseous reactants and products of chemical reactions are always in whole number ratios

Molar volume of gases
The volume that one mole of a gas occupies at a specified temperature and pressure V=molar vol.

Gas stoichiometry
Procedure:
1. Write a balanced chemical equation with given information written under substances
2. Find moles of known and use to find unknown moles using unknown/known ratio
3. Convert to appropriate final quantity

Not sure if vapour pressure is gonna be on the exam but just in case…
Vapor pressure
Water evaporates relatively easily and the gas collected will be mixed with some water vapor.
Water vapor is a gas like any other gas and the pressure exerted by a gas above its liquid is called VAPOR PRESSURE
As temperature increases, vapour pressure increases

Vapour pressure and Dalton’s Law
Dalton’s law of partial pressures and a table of known vapor pressures of water can be used to determine the pressure of dry gas that has been collected.

Dry gas is the natural gas without water vapour.