## Week 1 : Introduction, Chapters 1-4

### Mon, 4-Jan-16

Welcome to CHEM 120! Overview of syllabus and administrative matters. Tutorials and on-line MapleTA assignments begin next week.

### Wed, 6-Jan-16

Course notes, Module 1. Stoichiometry: mass vs. moles, molar mass; combining ratios, stoichiometric coefficients, and equivalents.

### Fri, 8-Jan-16

More stoichiometry: equivalents in stoichiometric calculations; reaction yield; simultaneous reactions, simultaneous equations; limiting reagents. Note: when performing calculations, use all digits; do not round to sig figs until the final result!

## Week 2 : Chapters 3-5

### Mon, 11-Jan-16

Limiting reagent and simultaneous reaction problems (course notes, pp. 1.9 and 1.10). Course notes, Module 2. Solutions; ionic vs. molecular compounds; complete vs. incomplete reactions.

### Wed, 13-Jan-16

Solutions; strong vs. weak electrolytes; concentration vs. density. Precipitation reactions, solubility rules; net ionic equations -- only dissociate strong electrolytes!.

### Fri, 15-Jan-16

Constructing and balancing molecular and net ionic equations for aqueous precipitation reactions.

## Week 3 : Chapters 4 and 5

### Mon, 18-Jan-16

Acid-base reactions; strong acids or bases = strong electrolytes, weak acids or bases = weak electrolytes; common strong and weak acids and bases. Oxidation state rules.

### Wed, 20-Jan-16

Oxidation states from electronegativities; balancing aqueous redox reactions.

### Fri, 22-Jan-16

Balancing aqueous redox reactions: syllabus and textbook methods. Course notes, Module 3. State variables for gases: P, V, and T; units, always use Kelvin for temperature!

## Week 4 : Chapter 6

### Mon, 25-Jan-16

Hydrostatic manometry. The ideal gas law and some applications.

### Wed, 27-Jan-16

The ideal gas law in chemical reactions. Gas mixtures: partial pressures and partial volumes.

### Fri, 29-Jan-16

Problem 3.7 solvable using volumes alone: Law of Combining Volumes. Partial pressures: gas trapping over water, vapor pressure of water. Kinetic theory of gases: kinetic energy, velocity, effusion and diffusion rates (velocities) and times vs. molar mass.

## Week 5 : Chapters 6-7 and Term Test 1

### Mon, 1-Feb-16

Guidelines and procedures for first term test. This information has been posted under the exam link. More on effusion and diffusion. Real gases.

### Wed, 3-Feb-16

Term Test 1.

### Fri, 5-Feb-16

Real gases: deviations from ideal gas behavior, causes and effects; van der Waals equation and virial equation of state. Course notes, Module 4. Introduction to thermodynamics: definitions -- systems, energy, heat (note sign convention!).

## Week 6 : Chapter 7

### Mon, 8-Feb-16

Heats of fusion and vaporization; heat capacities; "PV" work. Note sign conventions! "PV" work: work done by gas expansion and effect on internal energy and heat.

### Wed, 10-Feb-16

Internal energy. The first law of thermodynamics: heat and work. Enthalpy: internal energy plus the work needed to accommodate ("make space for") the system. Change in enthalpy equivalent to the heat added or released in the system under constant pressure. Enthalpy is useful to chemists because most reactions are conducted under constant pressure conditions and enthalpy gives total energy changes.

### Fri, 12-Feb-16

Internal energy, enthalpy, and PV: at small changes in volume (constant moles of gas), enthalpy changes approximately equal internal energy changes. Enthalpy changes for physical changes and chemical reactions. Enthalpies of formation: definition, reference states, use in calculation of reaction enthalpies. Endo- vs exothermicity.

## Week 7 : Chapter 8

### Mon, 22-Feb-16

Hess's law: adding reactions adds reaction enthalpies. Note that changes in enthalpy are the same regardless of the path taken, i.e., regardless of individual reaction steps; reaction enthalpies can be calculated even if the individual reactions used in the calculation do not occur for the net reaction. Course notes, Module 5. The nature of light: fundamental relationships, the electromagnetic spectrum. Note: the speed of light in vacuum (c) is an upper limit and is independent of frequency (i.e., is a "constant"); when light travels through transparent matter, its speed is slower and will depend on frequency.

### Wed, 24-Feb-16

The limits of classical physics: 1) Black body radiation and the UV catastrophe, quantization of energy and Planck's Law. 2) Photoelectric effect: particle nature of light. 3) Atomic emission and absorption (line) spectra: the Bohr model.

### Fri, 26-Feb-16

Bohr model: planetary model of electrons around nuclei; only certain orbits permitted, hence quantized energies; well-defined energy relationship for 1-electron atoms. Hydrogen emission series. Note: negative energy value = bound state; positive = unbound. De Broglie relationship: matter waves; wave-like behavior only observeable for very tiny masses; mathematically relates wave properties of matter to physical effects; interference (superposition, diffraction) is an intrinsic property of waves. Heisenberg uncertainty principle: measurement changes what's being measured; also a consequence of wave-particle duality.

## Week 8 : Chapter 8 (con't)

### Mon, 29-Feb-16

The Schrödinger equation; solution yields wavefunctions and associated energies. When applied to hydrogenic (i.e., 1-electron) atoms, obtain orbitals (hydrogenic electron wavefunctions) and new quantum numbers: n, l, ml; energy levels, shells, sub-shells as function of quantum numbers -- note that energy depends only on n for a hydrogenic atom.

### Wed, 2-Mar-16

Wavefunctions for 1 electron bound to nucleus = atomic orbitals. Orbitals analogous to 3-d standing waves. Meaning of ψ (the wavefunction), ψ2 (electron density, i.e., probability, at point). Nodes: regions where ψ = 0, where the electron cannot be; angular vs. radial nodes. Atomic orbitals (MUST KNOW!): effect of quantum numbers n, l, ml on orbital shape (s, p, d orbitals), nodes, and orientation, standard designations.

### Fri, 4-Mar-16

Radial component of the atomic orbital wave function: R (the radial wavefunction), R2 (electron density, i.e., probability, at point); 4πr2R2 (total probability of finding electron at all points at distance r); essential features of the radial component (nodes, extrema, size). Electron spin and ms.

## Week 9 : Term Test 2 and Chapters 8-9

### Mon, 7-Mar-16

Information for term test 2. Guidelines posted under exam link. Be sure to complete the identification section of the Scantron sheet correctly, especially student ID number and test version! Course Notes, Module 6. Multi-electron atoms: effect on shell and sub-shell energies; penetration and shielding in multi-electron atoms. The Aufbau Principle -- filling atomic energy levels to get the ground state configuration; Hund's Rule and the Pauli Principle. Writing quantum numbers to describe an electron in an orbital.

### Wed, 9-Mar-16

Term Test 2.

### Fri, 11-Mar-16

The periodic table: Z vs quantum numbers n and l; organization by groups, periods, blocks. Writing configurations to describe orbital populations; remember: nd < (n+1)s in metal cations! Electron pairing: diamagnetism vs. paramagnetism.

## Week 10 : Chapter 10

### Mon, 14-Mar-16

Periodicity of measurable properties: radii, ionization energies, electron affinities. Physical explanations for periodic trends and exceptions. For atom size, Zeff and principal quantum number. For IE and EA, Zeff (across a period), differing shell and sub-shell energies (filled shell stabilities), double orbital occupancy (half-filled sub-shell stability), spin alignment (Pauli "repulsion"); principal quantum number n (down a group). Beware of exceptions to general trends!

### Wed, 16-Mar-16

Course notes, Module 7. Nature of chemical bonds. Bonding limits: electron sharing = covalent, electron transfer = ionic. Octet rule.

### Fri, 18-Mar-16

Lewis structures. Procedure: 1) determine electron configurations and valence electron counts, 2) know or guess connectivity, 3) form initial connectivity (sigma bond) framework, 4) fill in remaining electrons as lone pairs or multiple bonds to achieve octets if possible. Dot and line notation. Be aware of multiple bonds, molecular charges, formal charges resonance forms, isomers. Assessing stability of Lewis structures.

## Week 11 : Chapter 10 (con't)

### Mon, 21-Mar-16

Lewis structure examples. Introduction to electronegativity.

### Wed, 23-Mar-16

Exceptions to the octet rule: 1) odd-electron species (free radicals), 2) hypervalent species (octet expansion) for heavier (n > 2) elements due to larger atom size, 3) electron-deficient species (early main group elements, esp. boron). Partitioning shared electrons: formal charges (covalent division) and oxidation states (ionic division). Note that this method of calculating oxidation state is, in fact, the definition of oxidation state. Bond length, strength, order. Bond energies: approximate average energies, use in estimating heats of reaction.

### Mon, 28-Mar-16

Electronegativity: determination of oxidation states, dipole moments. Molecular shapes and the VSEPR model: electron pair domains avoid each other. General protocol: 1) determine number of electron pairs about stereocenter and 2) treat multiple bonds as single electron pair domain. Wedge-dash notation. Standard VSEPR geometries, names, bond angles.

## Week 12 : Chapter 11

### Wed, 30-Mar-16

The VSEPR model: electron pair domains avoid each other, repulsions order as lp/lp > lp/bp > bp/bp. Molecular dipole moments: Lewis structures to VSEPR geometries to local dipoles (from electronegativies) vector summed to give molecular (net) dipole moment. Course notes, Module 8. Quantum mechanics, the electron pair (2-center/2-electron) bond, and valence bond theory: overlapping atomic orbitals produces bonds; σ bonds (direct overlap, no nodes coplanar with bond axis) vs. π bonds (indirect overlap, one node coplanar with bond axis).

### Mon, 1-Apr-16

Valence bond theory: hybridization gives bonding orbitals with proper directions; hybrid orbitals = linear combinations (sums and differences) of standard atomic orbitals; standard notation for hybrids and relevant geometries: sp = linear, sp2 = trigonal planar, sp3 = tetrahedral, sp3d = trigonal bipyramidal, sp3d2 = octahedral; hybridization and VSEPR electron counts; remaining atomic orbitals for multiple (π) bonds; in VB theory, an orbital (hybrid or atomic) can only be used to form one bond.

### Mon, 4-Apr-16

Molecular orbital theory: add and subtract atomic orbitals on different atoms to obtain bonding and antibonding combinations; phases are important! Combine s and p orbitals to get σ bonds; combine p orbitals to get π bonds. Bonding = extra electron density between nuclei; antibonding = less electron density (a node) between nuclei, i.e., perpendicular to the bond axis. Notation: σ = no node along bond axis; π = one node along bond axis; * = antibonding. Energy level diagram: per atomic orbitals, populate according to Aufbau, Pauli, Hund's rules to get ground state configuration. Bond order. MO energy level diagrams for light (2p) element diatomics; note -- σ(2p) and π(2p) level crossing between N2 and O2.