Review for third exam:
Chapter 6 (6.4 to 6.9)
De Broglie wavelength; electron diffraction
Heisenberg Uncertainty Principle ( x) (p)  h/4
Quantum numbers for the electron in hydrogen; restrictions and meaning
n - principle quantum number; relationship to energy and distance from nucleus
l - angular momentum quantum number; relationship to orbital type (s, p, d, ...)
shapes of s, p and d orbitals
nodes; relationship between n, , and number of radial nodes
ml - magnetic quantum number; relationship to the number of different orbitals
ms - spin quantum number; relationship to electron spin
Determining possible values for n, l, m, and ms
n = 1, 2, 3, …
Examples:
If n = 3, l = 0, 1, 2
l = 0, 1, 2, … (n-1)
If l = 2, ml = -2, -1, 0, +1, +2
ml = -l, …, +l
ms = - 1/2, +1/2
Electron configuration for atoms
Pauli principle
Aufbau principle
Hund's rule
Order of energy for orbitals 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < …; mnemonic device
Shorthand notation for electron configuration (using noble gases)
Orbital filling diagrams; counting unpaired electron spins; diamagnetic and paramagnetic
Electron configuration and the periodic table; core and valence electrons
Anomolous electron configurations; transfer of one s electron to give a half-filled or filled d
orbital (Cr, Mo, Cu, and Ag)
Chapter 7
Formation of ions by main group elements; relationship to noble gas configurations
Common ions for main group elements
Method for finding electron configurations for metal cations (write configuration for the atom,
then remove electrons from the highest n, or highest l (for orbitals with same n) to get correct charge)
Effective nuclear charge, Zeff; its calculation and interpretation
Trends in atomic size (atoms in the same group, atoms in the same row); explanation for trends
Trends in ion size (ions with the same charge in the same group, different ions of the same
element, ions with the same number of electrons); explanation for trends
Definition of first ionization energy and higher ionization energies
Trends in first ionization energy (atoms in the same group, atoms in the same row); explanation
for trends
Jumps in higher ionization energies; relationship to the number of valence electrons
Electron affinity; definition; relative values for halogens and noble gases
Making predictions about sizes of atoms and ions, ionization energies, and similar questions
Metallic character and its trends
Chapter 8
General types of bonding; ionic, covalent, metallic
General properties of ionic bonding; formation of binary ionic compounds
Lewis (dot) structures for atoms; dot structures for main group atoms; octet rule
Lewis picture of formation of ionic compounds; dot structure for ions
Lattice energy; definition; Born-Haber cycle (you do not have to memorize the cycle)
Trends in lattice energy (ion size, ion charge); explanation for trends
Failure of ionic bonding of nonmetals with nonmetals
Covalent bonding; bonding electrons; lone pair electrons
Bond order; single bond, double bond, triple bond; bond order vs bond length
Lewis structures for polyatomic molecules and ions
Comparison of ionic and covalent bonding and the effects of bonding on properties
Electronegativity; trends in electronegativity
Bond polarity; representation of polar bonds by partial charges or arrows
Nonpolar covalent, polar covalent, and ionic bonding and the relationship of these
to electronegativity differences in bonded atoms
Dipole moment as a measure of the polarity of a molecule
General methods for finding Lewis structures for molecules or ions obeying the octet rule
Covalent bonding and Lewis structures for organic molecules
Resonance structures
Formal charge and rule for assigning formal charge
Use of formal charge to find the "best" Lewis structure for a molecule or ion
Coordinate covalent bonds
Exceptions to the octet rule
Less than an octet of electrons (Be, B, Al)
Odd number of electrons (example: NO2)
Augmented (expanded) octets for elements in the third row and below (example: SF6)
Bond dissociation enthalpy; average bond dissociation enthalpy; trends in relation to bond order
Estimating Hrxn using a table of average bond energies
Average bond length; trends in relation to bond order
Metallic bonds and its general description; relationship to properties of metals
Chapter 9
Electron geometry and molecular geometry; definition; difference between the two
VSEPR (valence shell electron pair repulsion) theory; counting electron containing regions
Electron and molecular geometries and bond angles for the common cases (table in powerpoint)
2 regions (linear)
3 regions (trigonal planar; trigonal planar or nonlinear molecular)
4 regions (tetrahedral; tetrahedral, trigonal pyramid, nonlinear molecular)
5 regions (trigonal bipyramid; trigonal bipyramid, see saw, T-shape, linear molecular)
6 regions (octahedral; octahedral, square pyramid, square planar molecular)
Geometries for interior atoms in large molecules
Deviations from "pure" geometries
Drawing three dimensional structures for molecules and ions
Polar molecules and the relationship to polar bonds in the molecules
Valence bond theory - Overlap of atomic orbitals to make covalent bonds
Hybrid orbitals; relation between the number of electron containing regions and hybridization
(2 regions = sp, 3 regions = sp 2, 4 regions = sp3, 5 regions = sp3d, 6 regions = sp3d2)
Sigma () bond, pi () bonds - appearance, formation, counting sigma and pi bonds
Molecular orbital theory - LCAO-MO (linear combination of atomic orbitals to form molecular
orbitals)
Formation of  (bonding) and * (antibonding) molecular orbitals
Formation of  (bonding) and * (antibonding) molecular orbitals
Use of orbital filling diagrams for homonuclear diatomic molecules and ions
Electron configuration for homonuclear diatomic molecules and atoms
Bond order and its calculation in MO theory
Counting unpaired electrons; diamagnetic and paramagnetic molecules and ions
Multicentered pi-bonding; relationship to resonance structures
Pi-bonding in benzene
Chapter 10
General properties of gases
Pressure; definition, units for pressure, conversion between units
Barometer; manometer
Boyle's law
Charles' law
Relationships derived from Boyle's law and Charles' law
Avogadro's hypothesis
The ideal gas law
State variables; extensive and intensive variables; molar volume (V m = V/n)
Assumptions involved in the ideal gas law; limiting conditions for ideal behavior
The gas constant (R) in various units
Calculations involving the ideal gas law; determination of molecular mass from gas density
Gas density; STP (standard temperature and pressure)
Dalton's law of partial pressures; mole fraction
Kinetic theory; assumptions use in kinetic theory
rms average speed, kinetic energy per mole of gas
Maxwell velocity distribution
Diffusion, effusion; depenence on 1/M1/2
Real gases
van der Waals equation; interpretation of a an b coefficients; calculations with the equation