General Chemistry 120: Comprehensive Study Notes and Exam Review

General Chemistry Concepts and Problem Solving

Introduction

This study guide covers essential topics in General Chemistry, as reflected in a comprehensive final exam. The material spans atomic structure, chemical bonding, thermochemistry, gases, solutions, chemical kinetics, equilibrium, acids and bases, and more. Key equations and a periodic table are included for reference.

Atomic Structure and Periodic Properties

Electronegativity and Bond Polarity

• Electronegativity is the ability of an atom to attract shared electrons in a chemical bond.

• Differences in electronegativity between atoms determine the polarity of a bond.

• If the difference is large, the bond is polar; if small or zero, the bond is nonpolar.

• Example: In HCl, Cl is more electronegative than H, so the bond is polar.

Metathesis (Double Displacement) Reactions

• Metathesis reactions involve the exchange of ions between two compounds to form new compounds.

• General form: AB + CD → AD + CB

• Examples:

  • NaCl(aq) + AgNO3(aq) → NaNO3(aq) + AgCl(s)

  • BaCl2(aq) + Na2SO4(aq) → BaSO4(s) + 2NaCl(aq)

Balancing Chemical Equations

• When balancing equations, the number of atoms of each element must be the same on both sides.

• Subscripts in chemical formulas should not be changed; only coefficients are adjusted.

Formal Charge

• Formal charge is a bookkeeping tool to estimate the charge distribution in a molecule.

• Formula: $\text{Formal charge} = \text{Valence electrons} - \text{Nonbonding electrons} - \frac{1}{2}(\text{Bonding electrons})$

• Formal charges do not always represent the actual charge distribution.

Hybridization of Atomic Orbitals

• Hybridization explains the shapes of molecules by mixing atomic orbitals to form new hybrid orbitals.

• Types: sp, sp2, sp3, sp3d, sp3d2

• Hybrid orbitals differ in shape and energy from pure atomic orbitals.

Chemical Quantities and Stoichiometry

Mole Calculations

• The mole is a counting unit for atoms, molecules, or ions.

• Avogadro's number: $6.022 \times 10^{23}$ particles/mol

• Example: To find moles in a given mass: $\text{moles} = \frac{\text{mass (g)}}{\text{molar mass (g/mol)}}$

Empirical and Molecular Formulas

• The empirical formula gives the simplest whole-number ratio of atoms in a compound.

• The molecular formula gives the actual number of atoms of each element in a molecule.

• To determine empirical formula: Convert mass % to moles, divide by smallest, and write the ratio.

States of Matter and Intermolecular Forces

Critical Pressure and Supercritical Fluids

• The critical pressure is the minimum pressure required to liquefy a gas at its critical temperature.

• Above the critical temperature and pressure, a substance exists as a supercritical fluid.

Intermolecular Forces

• Types include London dispersion forces, dipole-dipole interactions, and hydrogen bonding.

• These forces affect boiling points, melting points, and solubility.

• Example: H2O exhibits hydrogen bonding, CH4 exhibits only London dispersion forces.

Electrolytes

• Strong electrolytes dissociate completely in solution (e.g., NaCl).

• Weak electrolytes dissociate partially (e.g., acetic acid).

Gases and Gas Laws

Ideal Gas Law

• Relates pressure, volume, temperature, and moles of a gas.

• Equation: $PV = nRT$

• Where P = pressure, V = volume, n = moles, R = gas constant, T = temperature (K).

Effusion and Diffusion

• Graham's Law: The rate of effusion is inversely proportional to the square root of molar mass.

• Equation: $\text{Rate} \propto \frac{1}{\sqrt{M}}$

Chemical Bonding and Molecular Structure

Lewis Structures

• Show the arrangement of valence electrons in molecules.

• Formal charges and resonance structures can be indicated.

Bond Angles and Molecular Geometry

• Determined by the number of bonding and lone pairs around a central atom (VSEPR theory).

• Example: Tetrahedral geometry has bond angles of 109.5°.

Thermochemistry

Heat, Work, and Internal Energy

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

• Where $\Delta U$ is change in internal energy, q is heat, w is work.

Enthalpy and Calorimetry

• Enthalpy change ($\Delta H$): Heat absorbed or released at constant pressure.

• Calorimetry: $q = mC\Delta T$

• Standard enthalpy of formation and combustion can be calculated using tabulated values.

Chemical Kinetics and Equilibrium

Reaction Rates

• Rate laws express the relationship between concentration and rate.

• General rate law: $\text{rate} = k[A]^m[B]^n$

Equilibrium

• At equilibrium, the rates of forward and reverse reactions are equal.

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

Acids, Bases, and Aqueous Equilibria

pH and pOH

• pH: $\text{pH} = -\log[H^+]$

• pOH: $\text{pOH} = -\log[OH^-]$

• At 25°C, $\text{pH} + \text{pOH} = 14$

Redox Reactions and Electrochemistry

Balancing Redox Reactions

• Redox reactions involve the transfer of electrons.

• Half-reactions are balanced for mass and charge.

Electrochemical Cells

• Standard cell potential ($E^\circ_{cell}$) can be calculated from standard reduction potentials.

• Nernst Equation: $E = E^\circ - \frac{0.0592}{n} \log Q$

Mathematical Operations and Reference Tables

Key Equations

• See the provided equation sheet for formulas covering thermodynamics, kinetics, equilibrium, acid-base chemistry, and more.

• Examples include:

  • $q = mC\Delta T$ (heat transfer)

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

  • $K_p = K_c(RT)^{\Delta n}$ (gas equilibrium)

Periodic Table

• The periodic table provides atomic numbers, atomic masses, and element symbols.

• Useful for determining electron configurations, periodic trends, and chemical reactivity.

Sample Reference Tables

Periodic Table of Elements

Group	1A	2A	3A	4A	5A	6A	7A	8A
Period 1	H							He
Period 2	Li	Be	B	C	N	O	F	Ne
Period 3	Na	Mg	Al	Si	P	S	Cl	Ar
Period 4	K	Ca	Sc	Ti	V	Cr	Mn	Fe

Common Equations

Equation	Description
$PV = nRT$	Ideal Gas Law
$q = mC\Delta T$	Heat transfer
$\Delta G = \Delta H - T\Delta S$	Gibbs Free Energy
$\text{pH} = -\log[H^+]$	pH calculation
$K = \frac{[\text{products}]}{[\text{reactants}]}$	Equilibrium constant
$E = E^\circ - \frac{0.0592}{n} \log Q$	Nernst Equation

Additional info:

• Some questions reference specific calculations (e.g., density, gas laws, calorimetry) and require use of the provided equations and periodic table.

• Students should be familiar with laboratory techniques, such as balancing equations, drawing Lewis structures, and predicting reaction products.