  
Chapters 1 and 2

States of matter: solids, liquids, and gases

Classification of matter: matter, pure substances, compound vs. element, mixtures: homogeneous (solutions) vs. heterogeneous mixtures

Methods of separating mixtures: distillation, chromatography, crystallization, filtration; solubility differences

Density = mass/volume; intensive property (vs. extensive properties)

Metric system and measurement; uncertainty and significant figures; unit multipliers

Precision vs. accuracy

Atomic theory-John Dalton; law of constant composition; law of multiple proportions; law of conservation of energy

J.J. Thomson discovers the electron in 1897 (cathode rays); charge/mass ratio for electron found as 1.76 x 10
8
C/g

Robert Millikan’s oil drop experiment yields the charge on the electron = 1.6 x
10
-19
C; sets mass of electron at 9.1 x 10
-31
kg

Ernest Rutherford discovers the nucleus in 1911; alpha scattering experiment; nucleus is 10,000 times smaller than atom, contains 99.99% of atoms mass and all the positive charge

Atomic number; mass number; isotopes; average atomic mass; ions; oxidation number

Periodic table; periods; columns; periodicity-similar chemical properties yields families of elements; metals, nonmetals, metalloids

Empirical formulas; molecular formulas

Nomenclature; naming systems for compounds and ions
Chapters 3 and 4

Writing balanced chemical equations; subscripts vs. coefficients; conserving atoms during chemical changes is based on work of Antoine Lavoisier (law of conservation of mass)

Identifying types of chemical reactions: combination vs decomposition reactions; combustion reactions

Formula weights; molecular weights; molar masses; atomic mass units vs. grams

(1 gram = 6.02 x 10
23
Stoichiometry: grams
amu)

moles
 molecules

atoms

Determining empirical formulas and molecular formulas

Limiting reactant, theoretical yield, actual yield, percent yield problems

Solutions: molarity = moles per liter of solution; electrolytes, solubility, solutes, solvents, “like dissolves like” = polar solutes dissolve in polar solvents



Exchange reactions; precipitates, spectator ions, molecular equation, ionic equation, net ionic equation

Acid = proton donor or substance dissolving in water to produce H
+1
ions; base = proton acceptor or substance dissolving in water to produce hydroxide ions, OH
-1

Arrhenius neutralization reactions: acid + base = salt + water; H
+1
+ OH
-1
= H
2
O

Oxidation-reduction reactions: oxidation = loss of electron(s); reduction = gain of electron(s); oxidizing agent; activity series; displacement (exchange) reaction:
(lab) Cu
(s)
+ 2Ag
+1
(aq)
= 2 Ag
(s)
+ Cu
+2
(aq)
; identifying oxidation numbers of monatomic ions and elements in a polyatomic ion
Chapters 6 and 7
E n

 13.6 eV n 2
E photon
 h  ; h = 4.15
 10 -15 eV  s = 6.63
 10 -34 J  s
E photon
 = h

1240 eV  nm
 mv

Electromagnetic waves have wavelength,

, frequency,

, and c =

= 3 x 10 8 m/s

The electronic structure of atoms has been worked out over time by putting our ideas to the test: Max Planck (quantum theory); Niels Bohr (solar system model);
Louis deBroglie (matter waves); Werner Heisenberg (uncertainty principle);
Erwin Schrodinger (wave equation) and Quantum Mechanics => quantum numbers and orbitals, energy states, valence electrons

Quantum of energy; photons; E = h

; line spectra and excitation vs. emission; hydrogen: E n
= -13.6 eV/n
2
, n = principal quantum number

Electrons have dualistic nature: sometimes behaving as particles and sometimes behaving as waves; deBroglie and

= h/mv; matter waves discovered in 1927 at
Bell telephone labs (Davisson and Germer)

Uncertainty principle: there is a limit as to how precise we can know the position and momentum of an electron in an atom: (
 x) x (
 mv) ≥ h/4π
 Schrodinger’s wave equation can be solved to yield position probabilities and energies of electrons in atoms

Quantum numbers: n, l
, and m l
identify distinct orbitals in atoms; hydrogen: for each level, n, there are n
2
orbitals; s, p, d, and f orbitals come in sets of 1, 3, 5, and
7
 n = 1, 2, 3, …; l
= 0, 1, …, n – 1; m l
= -
l to 0 to
– l

Pauli exclusion principle: no two electrons in an atom can have the same set of quantum numbers; magnetic spin quantum number, m s
, distinguishes between the two electrons that can coexist in an orbital

Auf bau principle: electrons are added to orbitals of increasing energy, in general, filling each to capacity before beginning the next; Hund’s rule: when electrons are

added to degenerate orbitals, they are added such that the electrons remain as unpaired as possible

Electron configurations: 1s
2
2s
2
2p
6 …etc; valence orbitals are generally the outer s and p orbitals; successive ionization energies provide evidence for energy levels and shells

Effective nuclear charge: the amount of positive charge in the nucleus that a particular electron in an atom can actually feel; screening effect of core electrons;
Z eff
= Z-S

Trends in atomic size, ionization energy, and electron affinity; metallic character trend
 Noble gases are the key to understanding the elements’ characteristics; unreactive gases have completed energy levels of electrons and a stable octet of valence electrons

Alkali metals, alkaline earth metals, halogens, noble gases: representative elements to be familiar with

Distinguishing metals and nonmetals: properties
Chapters 8 and 9

Chemical bonding: ionic, covalent, and metallic bonding

Ions of the representative elements have noble gas electron configurations; metals form positive ions (low ionization energies and low effective nuclear charge associated with valence electrons); nonmetals form negative ions (high ionization energies, more negative electron affinities, and higher effective nuclear charge associated with valence electrons)

Lattice energy and Born-Haber cycle describes the energetics associated with the formation of one mole of an ionic compound from its elements

Covalent bonding involves the sharing of electrons; half-filled orbitals overlap;
Lewis symbols and structures represent the bonding of nonmetals in molecular compounds

Octet rule: atoms tend to gain, lose, or share electrons in acquiring eight valence electrons

Multiple bonding results in stronger bonding and shorter bond distances

Bond polarity can be determine by electronegativity differences: ionic vs. polar covalent vs. nonpolar covalent bonding; electronegativity trends; bond dipoles

VSEPR theory: valence shell electron pair repulsion theory; electron domains position themselves so as to minimize repulsions

Electron domain geometries; linear, trigonal planar, tetrahedral, trigonal bipyramidal; and octahedral shapes

Molecular shapes are determined by the numbers of lone pairs (LP) and bond pairs (BP) of electrons: linear; trigonal planar and bent; tetrahedral and trigonal pyramidal and bent; trigonal bipyramidal and seesaw and T-shaped and linear; octahedral and square pyramidal and square planar

Valence bond theory: atoms bond covalently via orbital overlaps; sigma (

) bonds involve end-to-end overlapping of orbitals; pi (π) bonds form when unhybridized

p orbitals overlap sideways; sigma bonds have more overlap and are stronger than pi bonds

Hybridization of atomic orbitals: in order to explain the shapes of molecules, we have to hybridize atomic orbitals; s + p = two sp hybrids, directed at 180° with respect to each other; s + p + p = three sp
2
hybrids, directed at 120° with respect to each other; s + p + p + p = four sp
3
hybrids, directed at 109.5° with respect to each other; s + p + p + p + d = five sp 3 d hybrids, directed according to trigonal bipyramidal geometry; s + p + p + p + d + d = six sp
3 d
2
hybrids, directed according to octahedral geometry
 formal charge is the charge associated with an atom in a molecule = # valence electrons - # electrons assigned to atom in the molecule; formal charge can lead to a more stable Lewis structure when several alternatives are found

Resonance structures: equivalent Lewis structures that include multiple bonds; bond strengths and bond distances are averaged in resonance scenarios; delocalized electrons are shared between many electrons in resonance structures
Chapter 10

Atmospheric pressure: gas pressure = force/area; barometer, barometric pressure determined by the height of a column of mercury supported by the weight of the atmosphere

Manometers measure pressures of laboratory samples of gases; distinguish between open-end and closed-end manometers

Kinetic-molecular theory of gases: behavior of gases: gas laws; ideal gas is a hypothetical gas in which molecules take up no space and exert no electrical forces on each other; ideal gas law: PV = nRT

The kelvin temperature scale is the international metric unit of temperature; absolute zero of temperature = 0 K = -273 °C; K = °C + 273; at absolute zero, molecules have zero kinetic energy
 Boyle’s law: P ~ 1/V; P
1
V
1
= P
2
V
2
 Charles’ law: V ~ T; V
1
/T
1
= V
2
/T
2

Gay-Lussac’s law: P ~ T; P
1
/T
1
= P
2
/T
2
 Avogadro’s law: V ~ n; V
 Dalton’s law: P ~ n; P
1
1
/n
1
= V
2
/n
2
/n
1
= P
2
/n
2
; P total
= sum of partial pressures; P
1
=

1
P total where

1
is the mole fraction of component #1 in gas mixture

STP conditions = 273 K and 1.0 atm

Molar volume at STP = 22.4 L/mol (ideally)

Effusion vs. diffusion; effusion rates identify relative average molecular velocities of different gases; average molecular speeds are rms quantities ( r oot m ean s quare quantities); Graham’s law: the rms speed at a given temperature is inversely related to the square root of its molecular mass

The average translational kinetic energy of a collection of molecules is directly proportional to its Kelvin temperature

Tenets of the kinetic-molecular theory: molecules take up no space and exert no electrical forces on each other



Real gases do take up space and their molecules exert attractive forces on each other. Johannes van der Waals developed an equation that predicts the nonideal behavior of gases:
(P  an 2
V 2
)(V  nb)  nRT
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