Review for third exam

Chapter 6 (6.4 to 6.9)

De Broglie wavelength; electron diffraction

Heisenberg Uncertainty Principle ( x) (p)  h/4

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

m_l - magnetic quantum number; relationship to the number of different orbitals

m_s - spin quantum number; relationship to electron spin

Determining possible values for n, l, m_, and m_s

n = 1, 2, 3, … Examples: If n = 3, l = 0, 1, 2

l = 0, 1, 2, … (n-1) If l = 2, m_l = -2, -1, 0, +1, +2

m_l = -l, …, +l

m_s = - 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, Z_eff; 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

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 H_rxn 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

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 = sp², 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

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/M^1/2