General Chemistry II: Final Exam Study Guide & Formula Sheet

General Chemistry II: Final Exam Study Guide

Exam Structure and Coverage

This study guide summarizes the structure, topics, and key formulas relevant to a comprehensive General Chemistry II final examination. The exam is divided into three main sections: Multiple Choice, Visualizing Concepts, and Calculations. Each section assesses understanding of core chemical principles, problem-solving skills, and the ability to interpret and analyze chemical data and diagrams.

Exam Sections and Topics

Section I: Multiple Choice

• Gas Laws: Relationships between volume (V), pressure (P), temperature (T), and moles (n); ideal vs. non-ideal gas behavior; effusion and diffusion rates.

• Intermolecular Forces (IMF): Effects on boiling point, surface tension, capillary action, and structure.

• Thermochemistry: Calculating heat of vaporization using Clausius-Clapeyron equation, exothermic/endothermic and exergonic/endergonic processes, work done by/on a system, specific heat capacity calculations.

• Solid State Chemistry: Estimating number of atoms in unit cells (cubic, body-centered cubic, face-centered cubic), classification of alloys, semiconductor doping (n-type vs. p-type).

Section II: Visualizing Concepts

• Kinetics: Interpreting integrated rate law plots to determine rate laws, rate constants, and initial concentrations.

• Acid-Base Chemistry: Analyzing titration curves to determine pKa/Ka, acid strength, and concentration.

• Phase Changes: Reading phase diagrams to identify triple/critical points, states of matter, and effects of pressure/temperature changes.

• Thermochemistry/Phase Changes: Using heating curves to find heat of vaporization, specific heat capacity, melting/boiling points.

• Thermodynamics: Interpreting free energy profile diagrams to label ΔG, equilibrium, spontaneity, and exo/endothermicity.

• Electrochemistry: Labeling electrochemical cells, calculating emf, and determining electron flow direction.

• Gas Laws (Advanced): Calculating moles, mass, molar mass, and partial pressures using diagrams and equations.

Section III: Calculations

• Kinetics: Using initial rate tables to determine rate laws, rate constants, and rates for given concentrations.

• Chemical Equilibrium: Using ICE tables to find equilibrium concentrations and equilibrium constants (Kc).

• Thermochemistry: Calculating enthalpy changes (ΔHrxn) using enthalpies of formation, bond energies, and Hess's Law.

• Buffers: Writing acid/base reactions, identifying conjugate pairs, calculating buffer pH (Henderson-Hasselbalch equation), and determining concentrations.

• Electrochemistry: Writing balanced redox reactions, calculating standard emf, free energy, equilibrium constants, and non-standard emf using the Nernst equation.

Key Concepts and Definitions

Gas Laws

• Ideal Gas Law: $PV = nRT$

• Partial Pressure: $P = dRT$

• Effusion/Diffusion: Rate is inversely proportional to the square root of molar mass.

Thermochemistry

• Heat (q): $q = m s \Delta T$

• Enthalpy Change: $\Delta H = \Delta E + P \Delta V$

• First Law of Thermodynamics: $\Delta E = q + w$

• Clausius-Clapeyron Equation: $\ln P = -\Delta H_{vap}/RT + C$

Kinetics

• Rate Laws: $\text{rate} = k[A]^n$

• Integrated Rate Laws:

  • First Order: $\ln[A] = -kt + \ln[A]_0$

  • Second Order: $1/[A] = kt + 1/[A]_0$

  • Zero Order: $[A] = -kt + [A]_0$

• Arrhenius Equation: $k = Ae^{-E_a/RT}$

Chemical Equilibrium

• Equilibrium Constant (K): Relates concentrations of products and reactants at equilibrium.

• ICE Table: Used to calculate equilibrium concentrations.

Acid-Base Equilibria and Buffers

• pH and pOH: $\text{pH} = -\log[H^+]$

• Henderson-Hasselbalch Equation: $\text{pH} = \text{p}K_a + \log \frac{[A^-]}{[HA]}$

• Conjugate Acid-Base Pairs: Differ by one proton (H+).

Thermodynamics

• Gibbs Free Energy: $\Delta G = \Delta H - T \Delta S$

• Relationship to Equilibrium: $\Delta G = -RT \ln K$

• Spontaneity: $\Delta G < 0$ indicates a spontaneous process.

Electrochemistry

• Standard Cell Potential: $E^\circ_{cell} = E^\circ_{cathode} - E^\circ_{anode}$

• Relationship to Free Energy: $\Delta G^\circ = -nFE^\circ$

• Nernst Equation: $E = E^\circ - (RT/nF) \ln Q$

Solid State Chemistry

• Unit Cells: Simple cubic, body-centered cubic (BCC), face-centered cubic (FCC).

• Semiconductors: n-type (extra electrons), p-type (holes).

Formula Sheet

The following formulas and constants are provided for use during the exam:

• Gas Laws: $PV = nRT$, $P = dRT$

• Thermochemistry: $q = m s \Delta T$, $\Delta H = \Delta E + P \Delta V$, $\Delta E = q + w$

• Kinetics: $k = Ae^{-E_a/RT}$, $\ln[A] = -kt + \ln[A]_0$

• Equilibrium: $K = [\text{products}]/[\text{reactants}]$

• Acid-Base: $\text{pH} = \text{p}K_a + \log \frac{[A^-]}{[HA]}$

• Thermodynamics: $\Delta G = \Delta H - T \Delta S$, $\Delta G = -RT \ln K$

• Electrochemistry: $E = E^\circ - (RT/nF) \ln Q$, $\Delta G^\circ = -nFE^\circ$

• Constants: $R = 0.0821$ L·atm/mol·K or $8.314$ J/mol·K, $F = 96,485$ C/mol, $k$ (Boltzmann constant) $= 1.38 \times 10^{-23}$ J/K

Additional Information

• Be familiar with interpreting and constructing graphs, diagrams, and tables related to chemical concepts.

• Practice applying formulas to solve quantitative problems and explain conceptual relationships.

• Review the relationships between thermodynamic quantities, equilibrium, and electrochemical cell potentials.