History of the Periodic Table

Lesson 3.01: History of the Periodic Table
1. Explain how scientific observations led to the development of, and changes to, the periodic table.
Dmitri Mendeleev- first periodic table, organized 63 known elements according to properties, organized into rows and columns. He wrote names, mass and chemical properties on each.
Julius lothar Meyer- independently worked in German , similar to Mendeleev
Henry Gwyn Jeffrey’s Moseley: worked with Ernest Rutherford experimented with 38 metals, he found that the positive charge of each element nucleus increased by one from element to element as they were arranged in mendeleevs periodic table, lead to modern definition of atomic number (# of protons in atoms nucleus) and the recognition the atomic number was basis of organization of periodic table.
2. Describe the organization of the modern periodic table.
Arranged from left to right in rows (periods) by increasing atomic number and top to bottom in columns (groups) based on similar chemical properties.
Lesson 3.02: Group Names and Properties
1. Compare and contrast the properties of metals, metalloids, and nonmetals.
Metals: good conductors of heat and electricity and reflect light and heat, most luster 9shins) and most are malleable (hammered or rolled into sheets)
Non-metals: poor conductors of heat and electricity, most are gas at room temperature, those that are solid are not malleable
Metalloids: a semiconductor ( it conducts electricity better than non metals but not as good as metals) some characteristics of metals but more like non metals.
2. Identify groups and sections of the periodic table by group name and common properties.
Alkali Metals: in Group 1 of the periodic table are known as the alkali metals. These metals are extremely reactive, reacting with water and with most nonmetals. All of the alkali metals have a silvery appearance and they are soft enough to be cut with a knife. Alkali metals are so reactive that they do not occur in nature as pure elements; they are always found in compounds. Many of these compounds are essential to our lives, including being found in the fluids that fill and surround all of our cells. The metals of this group all have one valence electron, which they give up when they react to form ionic compounds.
Alkaline Earth Metals: The elements of Group 2 are called the alkaline-earth metals. These metals share similar properties because they each have two electrons in the s sublevel of their outermost energy level. Metals in Group 2 are harder, denser, and stronger than the metals in Group 1.Although they are less reactive than the Group 1 alkali metals, the alkaline-earth metals are still too reactive to be found as pure elements in nature. Compounds containing alkaline-earth metals are found in the earth’s crust, in sea salts, and in gems and minerals like emerald and aquamarine.

Hydrogen and Helium: Hydrogen is placed in Group 1 on the periodic table because its electron configuration is 1s1, but it is actually a nonmetal that does not resemble the alkali metals in this group.Helium has the electron configuration 1s2, meaning it has two valence electrons like the alkaline-earth metals in group 2. So, why is helium found in Group 18 with the noble gases? Like the other noble gases, helium has a full valence of electrons (remember that the first energy level only holds two electrons). There is a special stability related to a full valence, and helium shares this and other properties with the other nonmetals in Group 18.

Halogens: The elements in Group 17 are known as the halogens. The halogens are the most reactive of the nonmetals; they react with most metals to form ionic compounds known as salts. As pure elements, the halogens are diatomic molecules.This means that the halogens bond in pairs, forming compounds such as F2, Cl2, and Br2. As you will learn, the high reactivity of the halogens is based on the presence of seven valence electrons, one electron short of having a full valence. Some of the halogens (F2 and Cl2) are gases at room temperature, while others are liquids (Br2) or solids (I2). The halogens are used in water purification, photography, insecticides, bleaches, plastics, and many other practical applications.
Noble Gases: Group 18, the group at the far right of the periodic table, contains nonmetals called the noble gases. The first noble gas to be discovered on Earth was argon, which was discovered in 1894 by the English physicists John Williams Strutt and William Ramsay.Argon makes up one percent of Earth’s atmosphere, but it escaped notice for so long because it does not react with any other elements. Another noble gas, helium, was discovered earlier in the sun’s light spectrum, but Ramsay was the first to discover that helium is also found on Earth. The noble gases are not found in compounds in nature because their full valence energy levels lead to very low reactivity. All noble gases have a full valence; helium has two electrons in the first energy level, while all other noble gases have eight valence electrons.
Transition Metals: The metals in the d-block, groups 3–12, of the periodic table are called transition metals or transition elements. They are good conductors of electricity and are typically less reactive than the metals in groups 1 and 2. Many of these metals can be found as pure elements in nature, and are used in construction, electrical wiring, and even jewelry. You will find that you may be familiar with many of the transition metals, such as silver, gold, copper, iron, mercury, and platinum.

Lanthanides and Actinides: You have probably noticed that there is a section of the periodic table that is often cut out and shown underneath the rest of the table. The first row of this section is part of the sixth period, and represents the 4f sublevel. Because the f sublevel has seven orbitals, each able to hold two electrons, this section of the periodic table is 14 elements wide. The lanthanides, the 14 elements found between lanthanum (La) and hafnium (Hf) in Period 6, are all shiny and reactive metals. Chemists use modern separation techniques to isolate these elements to be used in practical applications, like inside television tubes. The 14 elements found between actinium (Ac) and rutherfordium (Rf) in Period 7 are called the actinides. The first four actinides can be found naturally on Earth, while the rest have only been synthesized in a laboratory. The actinides are unstable and radioactive, so these elements are not used very often in everyday chemistry.
Lesson 3.03: Periodic Trends
1. Describe and explain the trends for effective nuclear charge, atomic radius, ionic radius, and ionization energy across a period and down a group.
Effective nuclear charge: the charge (from nucleus) felt by the valence electrons after you have taken into account the number of shielding electrons that surround the nucleus
Atomic radius: half the distance between centers of two atoms of that element that are bonded together
Ionization energy: the energy required to remove on electron from an element, resulting in a positive ion.
Iconic radius: one half the diameter of an ion. (a positive ion is called a cation, a negative ion is called anion. Nonmetals usually become anions and metals usually become vation).
2. Predict the properties of an element based on the known patterns of the periodic table.

3. Describe and explain the periodic trends for electron affinity (honors).
Electron affinity: the energy involved when a neutral atom gains an electron
Becomes more negative for each element across a period from goup 1 to group 17 because of an increase in effective nuclear charge
Becomes less negative going down a group because each electron is being added to a higher energy level farther from the nucleus.
4. Explain the exceptions to the trend across a period for ionization energy (honors).
Noble gases : in group 18 all have positive electron affinity values. The noble gases must be forced to gain an electron because they already have a full valence energy level.
The alkaline earth metals in group 2 and the nonmetals in group 15 both have electron affinity values close to zero due to electron repulsion and effective nuclear charge.
Nitrogen, in group 15 does not form a stable -1 ion because when an additional. Electron is added to nitrogen’s valence level it is added to a 2p orbital that already has one electron. The weak attraction between the added electron and nitrogen’s nucleus iis why there is not much energy given off.

Lesson 3.04: Valence Electrons and Bonding
1. Define and compare ionic and covalent bonding.
Ionic bond: a chemical bond that results from electrostatic attraction between positive and negative ions, electrons are given up one by one atom and gained by another atom, and then those atoms are attracted to each other between a metal and nonmetals.
Covalent bonds: electrons are shared between two atoms neither atom completely gains or loses electrons between two nonmetals.
2. Relate your knowledge of the periodic trends to the chemical bonding exhibited by various elements.

Lesson 3.05: Ionic Bonding and Writing Formulas
1. Determine an element’s ionic charge based on its location on the periodic table.
Group1: 1+
Group2: 2+
Group3: 3+
Group4: 4+
Group5: 3-
Group6: 2-
Group7: 1-
Group 8: non reactive noble gases
2. Write the correct ionic formula when given two elements that bond ionically. remember the OIL RIG oxidation is the loss (-) reduction is the gain (+) when an element looses its electron it will have a -be sign, and when element gains it will have =ve chrge, in an ionic bond, those are represented through lewis structure, or with a plus/minus on the symbol of the element

Lesson 3.06: Covalent Bonding and Lewis Structures
1. Determine how many covalent bonds an atom needs in order to fill its valence shell, using the periodic table.
Must get to 8 valance electrons ex: group 17 needs one more valance electron group 6 needs 2 more valance electrons
2. Draw correct Lewis structures to model covalently bonded molecules when given the name or formula of the molecule. CH20
BF3
NCS-
3. Describe your observations and conclusions from the virtual lab.

Lesson 3.07: Intermolecular Forces
1. Use VSEPR theory to predict the shape of a molecule based on its Lewis structure.
The VSEPR theory is about geometry of compounds and electron locations. Valence shell electron pair repulsion (VSEPR)theory is a model used, in chemistry, to predict the geometry of individual molecules from the number of electron pairs surrounding their central atoms. It is also named Gillespie–Nyholm theory after its two main developers.
2. Compare and contrast intermolecular forces (London dispersion, dipole-dipole, hydrogen bonding, and ion-dipole).
London dispersion forces occur between all molecules and particles but are the only force of attraction between nonpolar molecules or noble has atoms. These forces are the weakest of the intermolecular forces. The London dispersion forces are caused by the motion of electron.
Dipole dipole forces: are electrostatic interactions of permanent dipoles in polar molecules the attractive forces that occur between the positive end of one polar molecule and the negative end another polar molecule tend to align the molecules to increase the attraction.
Hydrogen bonding: is a particular strong dipole dipole interaction in which hydrogen is covalently bonding to a highly electronegative element and attracted to the very electronegative element in another molecule. It occurs only in molecules containing N-H, O-H, or F-H bonds.
Ion dipole: forces are attractive forces that result from the electrostatic attraction between an ionic compound and a polar molecule. This interaction is most commonly found in solutions, especially in solutions of ionic compounds in polar solvents such as water.
3. Identify the intermolecular forces experienced by different compounds.
Intermolecular forces: the forces of attraction that occur between individual molecules.
Lesson 3.08: Naming Compounds
1. Correctly name covalent compounds, ionic compounds, and acids when given their formulas.
Mono-1, di-2, tri-3, tetra-4, penta-5, hexa-6, hepta-7 octa-8, nona-9, deca-10.
2. Write the formulas for ionic compounds, covalent compounds, and acids from their names.
3. Name hydrates or write the formula of a hydrate when given its name (honors).
Lesson 3.09: Molar Mass of Compounds
1. Calculate the molar mass of compounds from the formula. sum all the atomic mass of the constituent atoms. For example, the molar mass of NaCl can be calculated for finding the atomic mass of sodium (22.99 g/mol) and the atomic mass of chlorine (24.45 g/mol) and combining them.
2. Determine empirical formulas from percent by mass or mass data.

3. Determine the molecular formula from the empirical formula and molar mass of a substance.

4. Calculate the molar mass of a hydrate and determine the formula of a hydrate from experimental data (honors).

5. Determine the empirical formula of a compound from the mass of the products produced in experimental reactions (honors).

49.50 compound