Lecture contents

CH102 General Chemistry

Spring 2017

http://quantum.bu.edu/courses/ch102-spring-2017/contents.html
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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 20, 2017
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.
SOE: The role of AO symmetry, overlap, and relative energy. Only valence AOs form MOs; inner-shell AOs contribute negligibly to bonding. MO description of hydroxide, OH.
Lecture slides and lecture recording.

Lecture 2, Monday, January 23, 2017 Simple MO description of water, HOH, predicts a bond angle of 90o. 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.
Lecture slides and lecture recording.

Lecture 3, Wednesday, January 25, 2017 Calculating energy of hybrid orbitals. Hybrid orbital water AO-MO correlation diagram.
PDF notes: Polyatomic MO recipe.
σ and π bonding in formaldehyde, H2CO.
Lecture slides and lecture recording.

Lecture 4, Friday, January 27, 2017 Polyatomic MO recipe: Formic acid, HC(O)OH (localized π bonds). Polyatomic MO recipe: Formate, HC(O)O- (delocalized π bonds).
Lecture slides and lecture recording.

Lecture 5, Monday, January 30, 2017 Complete formate, HC(O)O- (delocalized π bonds).
Begin Mahaffy et al., chapter 11: States of matter. Macroscopic versus microscopic understanding of the ideal gas law.
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 6, Wednesday, February 1, 2017 Complete kinetic molecular theory. Root mean squared speed 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 7, Friday, February 3, 2017 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 is slightly larger than the most probable speed, 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 8, Monday, February 6, 2017 Units of pressure: Pascal (Pa), bar, and atm. Units of the gas constant: J versus L atm. Real gases: effect of intermolecular attractions present when gas particles encounter one another (van der Waals a). Real gases: effect of molecular size (van der Waals b).
Lecture slides and lecture recording.

Lecture 9, Wednesday, February 8, 2017 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. 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 are in equilibrium.
YouTube: Triple Point Demo, Tert Butyl Alcohol
Lecture slides and lecture recording.

Lecture 10, Friday, February 10, 2017 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.
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 they no longer interact with one another electrically.
YouTube: Supercritical transition of liquid Cl2
Lecture slides and lecture recording.

Lecture 11, Wednesday, February 15, 2017 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 17, 2017
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. Knowing relative values of lattice enthalpies and of aquation enthalpies is not sufficient to know relative values of enthalpies of solution. 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.
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.

Lecture 13, Tuesday, February 21, 2017 The value of the reaction quotient, Q, relative to its value at equilibrium, the equilibrium constant K, tell whether a reaction is spontaneous (Q < K), non-spontaneous (Q > K), or at equilibrium (Q = K). Therefore, knowing Q and K, we know what will happen to the relative amounts of reactants and products as time passes. Each chemical equation has its own equilibrium constant. However, the numerical value of an equilibrium constant changes depending on how the chemical equation is written, according the three rules: (1) Multiplying each stoichiometric coefficient by a constant a (which can be fractional) raises the equilibrium constant to that power. (2) Reversing a chemical equation changes its equilibrium constant to its reciprocal. (3) A chemical equation that is the sum of tow other chemical equations has an equilibrium constant equal to the product (not sum) of the equilibrium constant of the two other chemical equations.
Lecture slides and lecture recording.

Lecture 14, Wednesday, February 22, 2017 Equilibrium amounts (molarities, pressure, etc.) are different for different starting conditions. Given starting conditions, (1) evaluate the reaction quotient Q, (2) compare Q to the equilibrium constant K, and (3) based on whether Q is smaller or larger than K, shift reactants to products or shift products to reactants, respectively. If equilibrium is disturbed, by abruptly changing the amount of a reactant or product, once equilibrium is restored, the new equilbrium amounts will only partially have offset the change (Le Chatelier's principle)
Lecture slides and lecture recording.

Lecture 15, Friday, February 24, 2017
Begin Mahaffy et al., chapter 14: Acid-base equilibria in aqueous solutions. Water reacts with itself (to a very small extent) to form equal amounts of H3O+(aq) and OH-(aq). The higher the temperature, the more H3O+(aq) and OH-(aq) are formed. This means the pH = − log[H3O+] of pure water is different at different temperatures. At 25 oC, the pH of pure water is 7.0. At 50 oC, the pH of pure water is 6.6. This means the Kw = [H3O+] [OH-] is (10-7.0)2 at 25 oC and (10-6.6)2 at 50 oC. In both cases water is "neutral", because it contains equal amounts of H3O+(aq) and OH-(aq). In general if K increases with temperature, the chemical equilibrium is endothermic and if K decreases with temperature, the chemical equilibrium is exothermic.
Lecture slides and lecture recording.

Lecture 16, Monday, February 27, 2017 An acid in aqueous solution is anything that when added to water results in more [H3O+] than [OH-]. A strong acid reacts essentially 100%, and so a ca molar solution results in [H3O+] = ca M. A weak acid reacts much less than 100%, and so a ca molar solution results in [H3O+] << ca M. From [H3O+] alone, we cannot tell if an acid is weak or strong, for must know also how much acid is added to water. A strong acid HA results in mostly conjugate base A- relative to unreacted acid HA. A weak acid HA results in very little conjugate base A- relative to unreacted acid HA. Therefore, it is the proportion of A- to HA that distinguishes strong from weak acids.
Lecture slides and lecture recording.

Lecture 17, Wednesday, March 1, 2017
Lecture slides and lecture recording.

Lecture 18, Friday, March 3, 2017
Lecture slides and lecture recording.