Chemistry II Final Exam Review: Thermochemistry, Kinetics, Equilibrium, and Acids & Bases

Thermochemistry

Key Concepts in Thermochemistry

Thermochemistry is the study of the energy and heat associated with chemical reactions and physical transformations. It focuses on the exchange of energy between a system and its surroundings during chemical processes.

• Kinetic Energy: The energy of motion. For molecules, this is related to their movement.

• Potential Energy: Stored energy due to position or composition.

• q: Symbol for heat exchanged in a process. Positive q means heat is absorbed (endothermic); negative q means heat is released (exothermic).

• Exothermic: A process that releases heat to the surroundings (q < 0).

• Endothermic: A process that absorbs heat from the surroundings (q > 0).

• Enthalpy (H): The heat content of a system at constant pressure.

• Specific Heat Capacity (s): The amount of heat required to raise the temperature of 1 gram of a substance by 1°C.

• ΔH: Change in enthalpy; heat absorbed or released at constant pressure.

Key Equations:

• Heat transfer: $q = m \times s \times \Delta T$

• Enthalpy change (reaction): $\Delta H = \sum n \Delta H_f \, (\text{products}) - \sum m \Delta H_f \, (\text{reactants})$

Example: Calculating heat evolved when water is formed: Use the enthalpy of formation and stoichiometry to find total heat released.

Chemical Thermodynamics

Gibbs Free Energy, Entropy, and Spontaneity

Chemical thermodynamics studies the direction and extent of chemical reactions. It introduces the concepts of entropy (ΔS), Gibbs free energy (ΔG), and spontaneity.

• ΔS (Entropy): Measure of disorder or randomness in a system.

• ΔG (Gibbs Free Energy): Determines spontaneity of a process at constant temperature and pressure.

• Spontaneity: A process is spontaneous if it occurs without external intervention (ΔG < 0).

Key Equations:

• Entropy change: $\Delta S = \sum n S_f \, (\text{products}) - \sum m S_f \, (\text{reactants})$

• Gibbs free energy: $\Delta G = \Delta H - T \Delta S$

Example: Calculating ΔG° at a given temperature using enthalpy and entropy values.

Chemical Kinetics

Rates of Reaction and Reaction Mechanisms

Chemical kinetics is the study of reaction rates and the steps by which reactions occur. It explains how different factors affect the speed of chemical reactions.

• Rate of Reaction: Change in concentration of a reactant or product per unit time.

• Factors Affecting Rate: Concentration, temperature, surface area, catalysts, and nature of reactants.

• Molecularity: Number of molecules involved in an elementary step (unimolecular, bimolecular, termolecular).

• Order of Reaction: Power to which the concentration of a reactant is raised in the rate law.

• Rate Law: Mathematical expression relating reaction rate to reactant concentrations.

• Integrated Rate Law: Relates concentrations of reactants to time.

• Rate-Determining Step: The slowest step in a reaction mechanism, which controls the overall rate.

• Activation Energy (Ea): Minimum energy required for a reaction to occur.

• Catalyst: Substance that increases reaction rate by lowering activation energy, without being consumed.

• Collision Theory: Reactants must collide with sufficient energy and proper orientation to react.

• Reaction Mechanism: Sequence of elementary steps that make up a complex reaction.

• Elementary Step: A single step in a reaction mechanism.

Key Equations and Concepts:

Order	Differential Rate Law	Integrated Rate Law	Half-Life	Straight-Line Plot	Slope
0	Rate = k	$[A]_t = -kt + [A]_0$	$[A]_0/2k$	[A] vs t	-k
1	Rate = k[A]	$\ln([A]_t/[A]_0) = -kt$	$0.693/k$	ln[A] vs t	-k
2	Rate = k[A]^2	$1/[A]_t = kt + 1/[A]_0$	$1/(k[A]_0)$	1/[A] vs t	k

Example: Determining rate law from experimental data, calculating rate constants, and using integrated rate laws to find concentrations or time intervals.

Chemical Equilibrium

Dynamic Equilibrium and the Law of Mass Action

Chemical equilibrium occurs when the rates of the forward and reverse reactions are equal, resulting in constant concentrations of reactants and products.

• Equilibrium Constant (K): Ratio of product concentrations to reactant concentrations at equilibrium, each raised to the power of their coefficients.

• Kc: Equilibrium constant in terms of concentration (mol/L).

• Kp: Equilibrium constant in terms of partial pressures (atm).

• Kw: Ion product constant for water ($K_w = [H_3O^+][OH^-] = 1 \times 10^{-14}$ at 25°C).

• Ka, Kb: Acid and base dissociation constants.

• Equilibrium Constant Expression: For aA + bB ⇋ cC + dD, $K_c = \frac{[C]^c[D]^d}{[A]^a[B]^b}$

• Le Chatelier’s Principle: If a system at equilibrium is disturbed, it will shift to counteract the disturbance.

Example: Calculating Kc from equilibrium concentrations or percent dissociation.

Acids, Bases, and Aqueous Equilibria

Definitions, Properties, and Calculations

Acids and bases are substances that can donate or accept protons or produce hydroxide ions in solution. Their behavior is described by several definitions and properties.

• Acids: Substances that donate protons (Brønsted-Lowry), accept electron pairs (Lewis), or produce H+ in water (Arrhenius).

• Bases: Substances that accept protons (Brønsted-Lowry), donate electron pairs (Lewis), or produce OH- in water (Arrhenius).

• Salts: Ionic compounds formed from acid-base neutralization.

• Hydronium Ion (H3O+): The form in which protons exist in aqueous solution.

• Hydroxide Ion (OH-): The characteristic ion of bases in aqueous solution.

• Hydrolysis: Reaction of ions with water to form acidic or basic solutions.

• Common Ion: An ion present from more than one source in solution, affecting equilibrium.

• Buffer: Solution that resists changes in pH upon addition of acid or base, typically containing a weak acid and its conjugate base.

• pH: Measure of acidity; $pH = -\log[H_3O^+]$

• pOH: Measure of basicity; $pOH = -\log[OH^-]$

• Titration: Analytical method to determine concentration by reacting with a standard solution.

• Indicator: Substance that changes color at a particular pH, used to detect endpoint in titrations.

• Equivalence Point: Point in titration where stoichiometric amounts of acid and base have reacted.

• End Point: Point in titration where the indicator changes color, signaling completion.

Key Equations:

• $K_w = [H_3O^+][OH^-] = 1 \times 10^{-14}$

• $pH = -\log[H_3O^+]$

• $[H_3O^+] = 10^{-pH}$

• $pOH = -\log[OH^-]$

• $[OH^-] = 10^{-pOH}$

• $pH + pOH = 14$

Example: Calculating pH, pOH, and concentrations in acid-base solutions; buffer calculations; titration curves.

Sample Problem Types and Solutions

Representative Calculations

• Thermochemistry: Calculating heat evolved or absorbed using enthalpy changes and stoichiometry.

• Specific Heat: Determining heat gained or lost and specific heat capacity using $q = m \times s \times \Delta T$.

• Hess’s Law: Using enthalpies of formation to find reaction enthalpy.

• Gibbs Free Energy: Calculating ΔG° using ΔH and ΔS at a given temperature.

• Kinetics: Calculating average rate, rate constants, and using integrated rate laws for concentration/time.

• Rate Law Determination: Using experimental data to find reaction order and rate constant.

• Reaction Mechanisms: Predicting rate law from proposed mechanisms and identifying the rate-determining step.

• Equilibrium: Calculating equilibrium constants from percent dissociation or concentrations.

Summary Table: Key Formulas and Constants

Concept	Formula
Heat transfer	$q = m \times s \times \Delta T$
Enthalpy change	$\Delta H = \sum n \Delta H_f \, (\text{products}) - \sum m \Delta H_f \, (\text{reactants})$
Entropy change	$\Delta S = \sum n S_f \, (\text{products}) - \sum m S_f \, (\text{reactants})$
Gibbs free energy	$\Delta G = \Delta H - T \Delta S$
Kw (water)	$K_w = [H_3O^+][OH^-] = 1 \times 10^{-14}$
pH	$pH = -\log[H_3O^+]$
pOH	$pOH = -\log[OH^-]$
Relationship	$pH + pOH = 14$

Additional info:

• For all calculations, ensure units are consistent (e.g., convert grams to moles where necessary).

• When using logarithms for pH/pOH, use base 10.

• For equilibrium and kinetics, always check if the reaction order or stoichiometry affects the calculation.