General Chemistry 101: Final Exam Study Guide

General Chemistry 101: Final Exam Study Guide

Introduction

This study guide summarizes key concepts, formulas, and problem types commonly encountered in a first-year General Chemistry course, as reflected in a comprehensive final exam. Topics include atomic structure, chemical reactions, thermochemistry, equilibrium, acids and bases, electrochemistry, and solution chemistry. The guide is organized by major themes and includes definitions, equations, and example applications.

Atomic Structure and Periodic Properties

Subatomic Particles and Isotopes

• Atoms are composed of protons, neutrons, and electrons.

• Atomic number (Z): Number of protons in the nucleus; defines the element.

• Mass number (A): Total number of protons and neutrons.

• Isotopes: Atoms of the same element with different numbers of neutrons.

Example Table: Subatomic Particles in Ions

Symbol	Protons	Neutrons	Electrons	Net Charge
Fe2+	26	30	24	+2
Zn	30	35	30	0
Al3+	13	14	10	+3

Chemical Nomenclature and Formulas

Writing Formulas for Acids and Ions

• Bromous acid: HBrO2

• Hypobromous ion: BrO-

• Bromous acid: HBrO2

• Perbromic acid: HBrO4

Chemical Reactions and Stoichiometry

Balancing Equations and Limiting Reactants

• Balance chemical equations to obey the law of conservation of mass.

• Identify limiting reactants by comparing mole ratios of reactants used and required.

• Calculate theoretical yield and excess reactant remaining.

Example: For the reaction 2 C6H12 + 5 O2 → 2 H2C6H8O6 + 2 H2O, use molar masses to determine limiting reagent and product yield.

Thermochemistry

Enthalpy, Entropy, and Gibbs Free Energy

• Enthalpy change (ΔH): Heat absorbed or released at constant pressure.

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

• Gibbs Free Energy (ΔG): Determines spontaneity of a reaction.

Key Equations:

• $\Delta G = \Delta H - T\Delta S$

• $\Delta G = -RT \ln K_{eq}$

Example: Calculate ΔH and ΔS for a reaction using standard enthalpies and entropies of formation, then determine spontaneity at a given temperature.

Chemical Equilibrium

Equilibrium Constants and Calculations

• Kc: Equilibrium constant in terms of concentration.

• Kp: Equilibrium constant in terms of partial pressure.

• Relationship: $K_p = K_c(RT)^{\Delta n}$, where Δn = moles of gas (products) - moles of gas (reactants).

Example: For the reaction SOCl2 (g) ⇌ SO2 (g) + Cl2 (g), calculate equilibrium composition and Keq given initial and equilibrium amounts.

Acids, Bases, and Buffers

pH, pOH, and Buffer Calculations

• pH: $\mathrm{pH} = -\log[\mathrm{H}^+]$

• pOH: $\mathrm{pOH} = -\log[\mathrm{OH}^-]$

• Relationship: $\mathrm{pH} + \mathrm{pOH} = 14$ at 25°C

• Buffer solutions: Resist changes in pH upon addition of small amounts of acid or base.

• Henderson-Hasselbalch equation: $\mathrm{pH} = \mathrm{p}K_a + \log\left(\frac{[\text{A}^-]}{[\text{HA}]}\right)$

Example: Calculate the pH of a buffer made from sodium propionate and propionic acid, and determine the effect of adding strong acid or base.

Electrolytes and Solution Properties

Strong vs. Weak Electrolytes

• Strong electrolytes: Completely dissociate in solution (e.g., NaCl, HCl).

• Weak electrolytes: Partially dissociate (e.g., CH3COOH, NH3).

Thermodynamics and Spontaneity

State Functions and Properties

• State function: Property dependent only on the current state, not the path (e.g., enthalpy, entropy, internal energy).

• Extensive property: Depends on the amount of substance (e.g., mass, volume).

• Intensive property: Independent of amount (e.g., temperature, density).

Redox Reactions and Balancing Equations

Oxidizing and Reducing Agents

• Oxidizing agent: Substance that gains electrons (is reduced).

• Reducing agent: Substance that loses electrons (is oxidized).

• Balance redox equations in acidic or basic solution using the ion-electron method.

Example: Balance the equation: $\mathrm{AsO}_2^- + \mathrm{NO}_3^- \rightarrow \mathrm{HAsO}_4^{2-} + \mathrm{NO}_2$ in basic solution and identify the oxidizing and reducing agents.

Gas Laws and Physical Properties

Gas Law Calculations

• Ideal Gas Law: $PV = nRT$

• Calculate partial pressures, mole fractions, and percent composition of gases.

Example: Calculate the percent of argon in a light bulb at 27°C given pressure and volume.

Constants and Useful Equations

• Gas constant, $R = 0.08206$ L·atm·mol-1·K-1

• Avogadro's number, $N_A = 6.022 \times 10^{23}$ mol-1

• Standard temperature and pressure (STP): 1 atm, 0°C

• 1 atm = 760 mm Hg

• 1 Calorie = 4.184 J

Common Equations:

• Quadratic formula: $x = \frac{-b \pm \sqrt{b^2 - 4ac}}{2a}$

• $K_p = K_c(RT)^{\Delta n}$

• $\Delta G = \Delta H - T\Delta S$

• $\Delta G = -RT \ln K_{eq}$

• $\Delta S_{\text{univ}} = \Delta S_{\text{sys}} + \Delta S_{\text{surr}}$

• $\Delta H_{\text{rxn}} = \sum n\Delta H_f^{\circ} (\text{products}) - \sum n\Delta H_f^{\circ} (\text{reactants})$

Sample Problem Types

• Calculate pH, pOH, and buffer capacity.

• Determine equilibrium concentrations and constants.

• Balance redox reactions and identify agents.

• Apply gas laws to real-world scenarios.

• Use stoichiometry to find limiting reactants and yields.

• Calculate thermodynamic quantities and predict spontaneity.

Conclusion

This guide covers the foundational concepts and problem-solving strategies essential for success in General Chemistry. Mastery of these topics will prepare students for exams and further study in chemistry and related sciences.