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Chemistry Study Guide
This is a guide – it DOES NOT contain everything that will be covered on the exam. Use your
notes and the presentations for additional information and concepts.
The Elements, Periodic Table and Electrons
I. The Formation of the Elements
A. Nucleosynthesis
1. H isotope – deuterium
a. Extreme temperatures and pressure/density
2. H  He C  fusion into final element  Fe
B. Star Life Cycle
1. Nebula
2. Sun/Star
a. Balance between fusion and gravity
3. Red Giant
4. White Dwarf
5. Supernova
6. Neutron Star and Black Hole
II. Organizing the Elements – History of the Table
A. Antoine Lavoisier
B. Johann Dobereiner and the Law of Triads
C. John Newlands and the Law of Octaves
1. Both laws fall apart due to exceptions
D. Kekule and Meyer
1. Valency
D. Mendeleev – organized by atomic weight
E. Moseley – organized by atomic number
F. Seaborg
1. Transuranium elements
2. Island of Stability video – PBS/NOVA
III. The Periodic Table
A. Periods are the rows and represent the energy level (n) of valence e-
Valence Energy Level (n) Valence Shapes
1
s
2
s, p
3
s, p, d
4
s, p, d, f
5
s, p, d, f
6
s, p, d, f
7
s, p, d, f
Max. Val. e2
2, 6 = 8
2, 6, 10 = 18
2, 6, 10, 14 = 32
2, 6, 10, 14 = 32
2, 6, 10, 14 = 32
2, 6, 10, 14 = 32
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Chemistry Study Guide
This is a guide – it DOES NOT contain everything that will be covered on the exam. Use your
notes and the presentations for additional information and concepts.
1.
Shapes
a. S is sphere
b. P is principle or a figure 8
c. D is diffuse – four leaf clove
d. F is fundamental or 3d six leaf clover
2. Pauli Exclusion Principle each having an opposite spin.
3. Hund’s Rule (Share the Cookie Rule).
4. Aufbau Principle – electrons must fill lowest energy levels first before
moving to next higher level
5. Bohr Model – There are 7 principle energy levels (n = 1 through 7) of
electrons; K, L, M, N.
a. Maximum electron capacity per energy level is 2n2
6. Orbitals (ml) - For any l there are 2I + 1 orbitals in that level; deals w/ angular
momentum along the x, y, and z axes and orientation of electron cloud
7. Sublevels / Orbital Type (l) - s (spherical); p (principle/perpendicular); d
(diffuse); f (fundamental); n – 1 = # sublevels
8. Spins (ms) - ↑ +1/2; ↓ - ½
9. Octect Rule - Completed valence electron energy level
B. Groups/Families are columns and have same # valence e- for A (Primary) elements
1. 10 elements have symbols that don’t match the first letters of name
2. Hg and Br only two natural liquids – all others solid or gas
C. Metals in A, and those in B (Secondary) are Transition Metals
1. Metals are to left of stair line (includes Al)
2. Lanthanides and Actinides are Transition Metals (Period 6 and 7)
D. Nonmetals – to right of stair line (Halogens and Noble Gases)
E. Metalloids – have side along stair line
1. Excludes Al at times
F. Groups and Characteristics – KNOW THESE!
G. Periodic Law
1. System of Classifying Elements
a. Done by type of valence orbital (s,p,d,f)
b. Total number of valence electrons
c. Atomic properties of the element
d. Atomic number
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Chemistry Study Guide
This is a guide – it DOES NOT contain everything that will be covered on the exam. Use your
notes and the presentations for additional information and concepts.
Group I A. Alkali Metals (Group 1)
- do not occur elementally in nature
- have one valence electron, s block
- have low ionization energies
- form colorless ions, each with a +1 charge
- are reactive metals obtained by reducing the +1 ions in their natural compounds
- are stored under kerosene or other hydrocarbon solvent since react with water vapor or oxygen in air
- form water-soluble bases
- are strong reducing agents
- are good conductors of electricity and heat
- are ductile, malleable, and soft enough to be cut with a knife
- have a silvery luster, low density, and a low melting point
II A. Alkaline Earth Metals (2)
- do not occur elementally in nature
- occur most commonly as carbonates, phosphates, silicates, and sulfates
- occur naturally as compounds that are either insoluble or only slightly soluble in water
- contain two valence electrons, s block
- tend to lose two electrons per atom, forming ions with a +2 charge
- are less reactive than alkali metals
- primarily form ionic compounds
- react with water to form bases and hydrogen gas
- are good conductors of heat and electricity
- are ductile and malleable
- have a silvery luster
- include the naturally radioactive element radium
The B Groups. Transition Metals (3-12)
- consist of metals in groups 3 through 12
- contain one or two valence electrons, in d or f blocks
- are usually harder and more brittle than metals in groups 1 and 2
- have higher melting and boiling points that metals in groups 1 and 2
- are good conductors of heat and electricity
- are malleable and ductile
- have a silvery luster, except copper and gold
- include radioactive elements 89 through 109
- include mercury, the only metal that is liquid at room temperature, and U #92, the last natural element
- have chemical properties that tend to differ from each other
- tend to have two or more common oxidation states
- often form colored compounds
- may form complex ions
III A. Boron Family (13)
- do not occur elementally in nature
- are scarce in nature (except aluminum, which is the most abundant metallic element)
- have three valence electrons in p block
- are metallic (except boron, which is a solid metalloid)
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Chemistry Study Guide
This is a guide – it DOES NOT contain everything that will be covered on the exam. Use your
notes and the presentations for additional information and concepts.
- are soft and have low melting points (except boron, which is hard and has a high melting point)
- are chemically reactive at moderate temperatures (except boron)
IV A. Carbon Family (14)
- includes a nonmetal (carbon), two metalloids (silicon and germanium) and two metals (tin and lead)
- vary greatly in both physical and chemical properties
- occur in nature in both combined and elemental forms
- have four valence electrons in p block
- are relatively unreactive
- tend to form covalent compounds (tin and lead also form ionic compounds)
V A. Nitrogen Family (15)
- consists of two nonmetals (nitrogen and phosphorus), two metalloids (arsenic and antimony), and one
metal (bismuth)
- nitrogen is most commonly found as atmospheric N2, phosphorus as phosphate rock, and arsenic,
antimony, and bismuth as sulfides or oxides; antimony and bismuth are also found elementally
- range from very abundant elements (nitrogen and phosphorus) to relatively rare elements (arsenic,
antimony, and bismuth)
- have five valence electrons in p block
- tend to form covalent compounds, most commonly with oxidation numbers of +3 or +5
- are solids at room temperature, except nitrogen
VI A. Oxygen Family (16)
- occur elementally in nature and in combined states
- consists of three nonmetals (oxygen, sulfur, and selenium), one metalloid (tellurium), and one metal
(polonium)
- have six valence electrons in p block
- tend to form covalent compounds with other elements
- tend to exist as diatomic and polyatomic molecules, such as O2, O3, S6, S8, and Se8
- commonly exist in compounds with the -2 oxidation state, but often exhibit other oxidation states
VII A. Halogen Family (17)
- are nonmetals and occur in combined form in nature, mainly as metal halides
- form salts when react with alkalines
- are found in the rocks of Earth's crust and dissolved in sea water
- range from fluorine, the 13th most abundant element, to astatine, which is one of the rarest
- exist at room temperature as a gas (F2 and Cl2), a liquid (Br2), and a solid (I2 and At)
- have seven valence electrons
- tend to gain one electron to form a halide, X- ion, also share electrons and have positive oxidation states
- are reactive, with fluorine being the most reactive of all nonmetals
VIII A. Noble Gases (18)
- not reactive
- have a full outer energy level, completed octet rule (p block)
- are all gases
- are all nonmetals
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Chemistry Study Guide
This is a guide – it DOES NOT contain everything that will be covered on the exam. Use your
notes and the presentations for additional information and concepts.
Hydrogen
- metal and nonmetal, not in any group
- proterium, deuterium and triterium are its isotopes
IV. The Electron
A. General Concepts
1. Electrons form when neutrons decay
2. Are leptons, non-baryonic matter
3. Make up beta radiation
4. Held to the nucleus through EMF
5. Normal orbit is called the ground state
a. If atom absorbs energy – electrons move up 1+ orbits to excited state
b. When energy lost – electrons return and release energy as photons
i. Each photon has different wavelength (color) depending on number of
orbits jumped and dropped
6. Are what determines if atom is an ion or not
a. If a loss of electrons – then + charge  cation
b. If a gain of electrons – then – charge  anion
B. Review
a. Thompson and Plum Pudding Model
b. Schrodinger and electron cloud
c. Heisenberg and Uncertainty Principle
d. De Broglie – electron wave function
e. Millikan – mass and charge of electron
C. Electron Configuration
1. Mapping valence electrons
a. For Primary elements, use column number, period and shape
b. For Secondary elements, can use period and shape, but must map out electrons to
determine valence numbers
2. Long and Short Form of Aufbau Diagrams
D. Electron Series
a. Wavelength and Frequency of photon
i. Lyman – change is 1 level; UV light
ii. Balmer – 2 levels; R,B,V
iii. Paschen – 3; infrared
iv. Brackett – 4 levels
v. Pfund – 5 levels
vi. Humphreys – 6 levels
V. Electron Configuration
A. Aufbau – Long and Short Form
B. Noble gas Configuration
C. Lewis Dot Diagrams