Lecture contents

Spring 2018

Department of Chemistry | Boston University

This page lists the contents of each lecture.

Use "Find" in your Web browser to find in this page those lectures where a particular topic is discussed.

For each lecture: there is a link to the PDF of the PowerPoint slides used in that lecture and a link to the lecture recording, showing the notations made to each slide during the lecture.

Lecture 1, Friday, January 19, 2018
Continue Mahaffy et al., chapter 10: Modeling bonding in molecules.
PDF notes: Bonding in diatomic molecules.
Which AOs form bonding and antibonding MOs? Inner-shell AOs overlap negligibly and so do not contribute to bonding.
CDF animation: Li2 1s and 2s bonding electron clouds.
CDF animation: Li2 1s and 2s antibonding electron clouds.
CDF animation: Li2 2s bonding and antibonding electron clouds.
CDF animation: XY 2s correlation diagram.
CDF animation: XY 2s correlation diagram, bonding electron cloud and its dipole moment.
CDF animation: XY 2s correlation diagram, antibonding electron cloud and its dipole moment.
SOE: The role of AO symmetry, overlap, and relative energy.
PDF notes: Questions on Symmetry, Overlap, Energy.
Only valence AOs form MOs; inner-shell AOs contribute negligibly to bonding or antibonding.
Lecture slides and lecture recording.

Lecture 2, Monday, January 22, 2018 Covalent versus ionic character in bonding. MO description of hydroxide, OH.
Lecture slides and lecture recording (video only).

Lecture 3, Wednesday, January 24, 2018 Simple MO description of water, HOH, predicts a bond angle of 90o.
PDF notes: Hybrid AOs and Polyatomic MOs.
Hybrid AOs (sp, sp2, and sp3) account for molecular shape.
CDF animation: Hybrid Orbitals in Organic Chemistry.
Hybrid orbitals have energies intermediate between those of the constituent AOs. Calculating energy of hybrid orbitals. Hybrid orbital water AO-MO correlation diagram.
Lecture slides and lecture recording.

Lecture 4, Friday, January 26, 2018
PDF notes: Polyatomic MO recipe.
$$\sigma$$ and $$\pi$$ bonding in formaldehyde, H2CO.
Formic acid, HC(O)OH (localized $$\pi$$ bonds).
Formate, HC(O)O- (delocalized $$\pi$$ bonds).
Lecture slides and lecture recording.

Lecture 5, Monday, January 29, 2018 Complete formate, HC(O)O- (delocalized $$\pi$$ bonds). Postscript on recipe for polyatomic MO recipe.
Begin Mahaffy et al., chapter 11: States of matter.
Macroscopic versus microscopic understanding of the ideal gas law.
Lecture slides and lecture recording.

Lecture 6, Wednesday, January 31, 2018 Review: Review: Localized and delocalized $$\pi$$ orbitals versus localized and delocalized electrons.
PDF notes: Kinetic molecular theory
Derivation of pressure in terms of time rate of change of momentum per collision with a container wall. Pressure is due to collisions of all $$N$$ particles of gas.
Lecture slides and lecture recording.

Lecture 7, Friday, February 2, 2018 Complete kinetic molecular theory. Root mean squared speed, $$u_{rms}$$ goes up with temperature and down with speed. Particle picture of gases. Understanding gas behavior in terms of motion and speed of individual particles. Mixtures of gases.
Lecture slides and lecture recording.

Lecture 8, Monday, February 5, 2018 Review: Practice with particle picture of gases. Units of pressure: Pascal (Pa), bar, and atm. Units of the gas constant, $$R$$: J versus L atm. Distribution of molecular speeds is due to collisions of gas particles with one another. After not too many collisions the relative number of particles with a given speed is given by the Maxwell-Boltzmann distribution. The higher the temperature, the broader but lower the distribution. The rms speed, $$u_{rms}$$, is slightly larger than the most probable speed, $$u_{mp} = \sqrt{2/3} u_{rms}$$, because of the exponential shape of the speed distribution at hight speeds. In calculating rms speed, it is crucial to carefully cancel units.
PDF article: Bonomo & Riggi, 1984, The evolution of the speed distribution for a two-dimensional ideal gas: A computer simulation.
Lecture slides and lecture recording.

Lecture 9, Wednesday, February 7, 2018 Complete molecular speeds and their distribution.
CDF animation: Maxwell-Boltzmann distribution of speeds.
Real gases: Van der Waals $$a$$ reflects intermolecular attractions present when gas particles encounter one another; therefore, hydrogen bonding can affect the value of van der Waals a. Because of the random orientation of close encounters, dipole-dipole interaction is typically less strong than dispersion interaction.
Lecture slides and lecture recording.

Lecture 10, Friday, February 9, 2018 Real gases: effect of molecular size (van der Waals $$b$$). Van der Waals equation. Conditions which require taking into account deviations of gas behavior from that described by the idea gas law, $$PV=nRT$$. Phase diagram lines are combinations of pressure and temperature at which two different phases (solid-liquid, liquid-gas, and solid-gas) are simultaneously present and in equilibrium. The triple point is when all three phases---solid, liquid, and gas---are in equilibrium.
YouTube: Triple point video, Tert Butyl Alcohol
The critical point is the temperature and pressure above which it is not meaningful to distinguish between liquid (too diffuse) and gas (too dense)---the supercritical region. By passing through the critical region, gas can be converted to liquid without any phase transition, that is, without gas and liquid being simultaneously present.
YouTube: Supercritical fluid video, Liquid Cl2
Lecture slides and lecture recording.

Lecture 11, Wednesday, February 14, 2018
Begin Mahaffy et al., chapter 12: Solutions and their behavior.
Ionic solids dissolve as a result of competition between attraction of oppositely charged ions in the solid and attraction of polar water molecules for the individual ions. Some ionic solids release heat when dissolving, and some absorb heat when dissolving.
Lattice enthalpy is the enthalpy change required to separate one mole of ionic solid into its individual ions in the gas phase, so that they are so far apart that they no longer interact with one another electrically. Lattice enthalpy is always positive, since energy is required to separate oppositely charged ions from one another.
Enthalpy of aquation (enthalpy of hydration) is the enthalpy change when one mole of oppositely charged ion pairs, initially in the gas phase, so that they are so far apart they no longer interact with one another electrically, is place in liquid water. Enthalpy of aquation is always negative, due to the attractive interaction of the polar water molecules with the individual ions.
Enthalpy change of solution is the enthalpy change when one mole of an ionic solid dissolves in water. By Hess's law, enthalpy of solution is the sum of the lattice enthalpy and the enthalpy of aquation. Since lattice enthalpy and the enthalpy of aquation are each large but opposite in sign, the magnitude of their sum is much smaller and whether enthalpy of solution is positive (endothermic) or negative (exothermic) cannot be predicted without further information.
Relative values of lattice enthalpy are determined by relative values of the Coulomb interaction energy between the oppositely charged ions. The smaller the ions, the closer they can be and so the larger the lattice enthalpy. The greater the charge on the ions, the larger the lattice enthalpy.
Lecture slides and lecture recording.

Lecture 12, Friday, February 16, 2018 Enthalpy of aquation (enthalpy of hydration) is the enthalpy change when one mole of oppositely charged ion pairs, initially in the gas phase, so that they are so far apart they no longer interact with one another electrically, is place in liquid water. Enthalpy of aquation is always negative, due to the attractive interaction of the polar water molecules with the individual ions.
Crystal ionic radii values: https://en.wikipedia.org/wiki/Ionic_radius
Relative values of aquation (hydration) enthalpy are determined by relative values of the Coulomb interaction energy between ions and surrounding polar water molecules. The smaller the ions, the closer the water molecules can be and so the larger the aquation (hydration) enthalpy. The greater the charge on the ions, the greater the attraction between the ions and the water molecules and so the larger the aquation (hydration) enthalpy. Enthalpy change of solution is the enthalpy change when one mole of an ionic solid dissolves in water. By Hess's law, enthalpy of solution is the sum of the lattice enthalpy and the enthalpy of aquation. Since lattice enthalpy and the enthalpy of aquation are each large but opposite in sign, the magnitude of their sum is much smaller and whether enthalpy of solution is positive (endothermic) or negative (exothermic) cannot be predicted without further information. Colligative properties are due to the presence of nonvolatile solute particles. There are four colligative properties: vapor pressure lowering, boiling point elevation, freezing point lowering, and osmotic pressure. At this point you are responsible for calculating colligative properties using provided formulas, in lab and in discussion. Later, we will learn where these formulas come from, taking into account the affect of the solute on the entropy of the system.
Lecture slides and lecture recording.

Lecture 13, Tuesday, February 20, 2018 Example osmotic pressure problem and answer.
Begin Mahaffy et al., chapter 13: Dynamic and chemical equilibrium.
If reactants are consumed to form products as time passes, a reaction is said to be spontaneous. If products are consumed to form reactants as time passes, a reaction is said to be non-spontaneous. If the amounts of reactants and products does not change with time, a reaction is said to be at equilibrium. Reaction quotient Q tell where amounts of reactants and products are relative to equilibrium.
Lecture slides and lecture recording.