General Chemistry II: Exam II Review Study Guide

Chapter 4: Chemical Reactions and Chemical Quantities

Writing Chemical Equations

Chemical equations represent chemical reactions, showing the reactants and products, their physical states, and the stoichiometric relationships between them.

• Balancing Chemical Equations: Ensure the number of atoms of each element is the same on both sides of the equation. This follows the Law of Conservation of Mass.

• Molecular Equations: Show all reactants and products as compounds, not ions.

• Full Ionic Equations: Represent all strong electrolytes as ions.

• Net Ionic Equations: Show only the species that undergo a chemical change, omitting spectator ions.

• Spectator Ions: Ions that do not participate in the actual chemical reaction.

Example: For the reaction of NaCl(aq) and AgNO3(aq):

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

• Full Ionic: Na+(aq) + Cl-(aq) + Ag+(aq) + NO3-(aq) → Na+(aq) + NO3-(aq) + AgCl(s)

• Net Ionic: Ag+(aq) + Cl-(aq) → AgCl(s)

Stoichiometry

Stoichiometry involves quantitative relationships between reactants and products in a chemical reaction.

• Reactant Calculations: Use mole ratios from the balanced equation to determine how much of one reactant reacts with another.

• Product Calculations: Calculate the amount of product formed from a given amount of reactant.

• Density as a Conversion Factor: Use density to convert between mass and volume: $\text{Density} = \frac{\text{Mass}}{\text{Volume}}$

Limiting Reactant Problems

The limiting reactant is the reactant that is completely consumed first, limiting the amount of product formed.

• Identifying Limiting Reactant: Compare the mole ratios of reactants to the balanced equation.

• Product Calculation: The amount of product is determined by the limiting reactant.

• Excess Reactant: Calculate the amount of reactant left over after the reaction.

Theoretical Yield and Percent Yield

• Theoretical Yield: The maximum amount of product that can be formed from the limiting reactant.

• Percent Yield: The ratio of actual yield to theoretical yield, expressed as a percentage: $\text{Percent Yield} = \frac{\text{Actual Yield}}{\text{Theoretical Yield}} \times 100\%$

Chapter 5: Introduction to Solutions and Aqueous Solutions

Solution Chemistry

Solutions are homogeneous mixtures of two or more substances. The concentration of a solution is often expressed as molarity.

• Molarity (M): $M = \frac{\text{moles of solute}}{\text{liters of solution}}$

• Using Molarity as a Conversion Factor: Molarity can convert between moles and volume.

• Dilution Calculations: $M_1V_1 = M_2V_2$ (where 1 = initial, 2 = final)

• Solution Stoichiometry: Use molarity and volume to determine moles for stoichiometric calculations.

Chemical Reactions in Solution

• Double Displacement Reactions: Two compounds exchange ions to form two new compounds.

• Solubility Rules: Used to predict whether a precipitate will form in double displacement reactions.

• Combustion Reactions: A substance reacts with O2 to produce CO2 and H2O (for hydrocarbons).

• Gas-Evolution Reactions: Reactions that produce a gas as a product.

• Redox Reactions: Involve the transfer of electrons between species.

• Oxidation: Loss of electrons; Reduction: Gain of electrons.

• Oxidizing Agent: Causes oxidation (is reduced); Reducing Agent: Causes reduction (is oxidized).

• Oxidation Numbers: Assigned to atoms to track electron transfer.

• Single Displacement Reactions: An element replaces another in a compound; activity series predicts feasibility.

• Acid/Base (Neutralization) Reactions: Acid reacts with base to form water and a salt.

Chapter 6: Gases

Gas Laws and Properties

Gases have unique properties and are described by several laws relating pressure, volume, temperature, and amount.

• Manometer: Device to measure gas pressure.

• Simple Gas Laws:

  • Boyle's Law: $P_1V_1 = P_2V_2$ (at constant T, n)

  • Charles's Law: $\frac{V_1}{T_1} = \frac{V_2}{T_2}$ (at constant P, n)

  • Avogadro's Law: $\frac{V_1}{n_1} = \frac{V_2}{n_2}$ (at constant P, T)

• Combined Gas Law: $\frac{P_1V_1}{T_1} = \frac{P_2V_2}{T_2}$

• Ideal Gas Law: $PV = nRT$

• Density of a Gas: $d = \frac{MP}{RT}$ (M = molar mass)

• Mass of a Gas: $n = \frac{m}{M}$

• Molar Mass of a Gas: $M = \frac{dRT}{P}$

• Dalton’s Law of Partial Pressures: $P_{\text{total}} = \sum P_i$

• Collecting Gases over Water: Account for vapor pressure of water: $P_{\text{gas}} = P_{\text{total}} - P_{\text{H}_2O}$

Gases in Chemical Reactions

• Stoichiometry with Gases: Use molar volume at STP (22.4 L/mol) for conversions.

• Limiting Reactant Problems: Apply gas laws to determine limiting reactant and product amounts.

Kinetic Molecular Theory and Gas Behavior

• Kinetic Molecular Theory: Explains gas properties based on particle motion.

• Temperature and Molecular Velocities: Higher temperature increases average kinetic energy and velocity.

• Diffusion and Effusion: Diffusion is mixing of gases; effusion is passage through a small opening.

• Graham’s Law of Effusion: $\frac{\text{Rate}_1}{\text{Rate}_2} = \sqrt{\frac{M_2}{M_1}}$

• Real Gases vs. Ideal Gases: Real gases deviate from ideal behavior at high pressure and low temperature due to intermolecular forces and finite molecular volume.

Chapter 8: The Quantum-Mechanical Model of the Atom

Wave Properties and Electromagnetic Spectrum

• Characteristics of a Wave: Wavelength (λ), frequency (ν), amplitude, speed (c).

• Electromagnetic Spectrum: Range of all types of electromagnetic radiation. $c = \nu \lambda$

Particle Nature of Light and the Photoelectric Effect

• Photoelectric Effect: Light can eject electrons from a metal surface; demonstrates light is quantized (photons).

• Energy of a Photon: $E = h\nu$ or $E = \frac{hc}{\lambda}$

• Wave-Particle Duality: Light and electrons exhibit both wave and particle properties.

Atomic Spectroscopy and the Bohr Model

• Emission and Absorption: Atoms absorb energy and electrons move to higher energy levels; emission occurs when electrons return to lower levels.

• Bohr Equation: $\Delta E = B \left( \frac{1}{n_f^2} - \frac{1}{n_i^2} \right)$

• Rydberg Equation: $\frac{1}{\lambda} = R \left( \frac{1}{n_1^2} - \frac{1}{n_2^2} \right)$

• Constants: $B$ and $R$ are experimentally determined constants.

Quantum Numbers and Atomic Orbitals

• Heisenberg’s Uncertainty Principle: Impossible to know both position and momentum of an electron simultaneously.

• Four Quantum Numbers:

  • Principal (n): Energy level (shell)

  • Angular Momentum (l): Sublevel (subshell)

  • Magnetic (ml): Orientation of orbital

  • Spin (ms): Electron spin direction (+1/2 or -1/2)

• Principal Level (Shell): Defined by n

• Sublevel (Subshell): Defined by l (s, p, d, f)

• Number of Subshells per Shell: Equal to n

• Number of Orbitals per Subshell: $2l + 1$

• Number of Electrons per Subshell: $2(2l + 1)$

• Atomic Orbitals: Regions in space where electrons are likely to be found; shapes depend on quantum numbers.

• Pauli Exclusion Principle: No two electrons in an atom can have the same set of four quantum numbers.

Chapter 9: Periodic Properties of the Elements

Electron Configuration and Periodic Trends

• Core and Valence Electrons: Core electrons are inner electrons; valence electrons are in the outermost shell and determine chemical properties.

• Electron Configurations: Show the arrangement of electrons in an atom or ion; anomalies exist for some transition metals.

• Orbital Diagrams: Visual representations of electron configurations using boxes and arrows.

• Paramagnetic: Atoms with unpaired electrons; Diamagnetic: All electrons are paired.

• Orbital Blocks: s, p, d, f blocks on the periodic table correspond to sublevel filling.

• Hund’s Rule: Electrons fill degenerate orbitals singly before pairing.

• Pauli Exclusion Principle: No two electrons in the same atom can have identical quantum numbers.

• Transition Elements: d-block elements with unique electron configurations.

Effective Nuclear Charge and Coulomb’s Law

• Effective Nuclear Charge (Zeff): The net positive charge experienced by valence electrons; increases across a period.

• Coulomb’s Law: $E = k \frac{q_1 q_2}{r}$ (describes the force between charged particles)

Isoelectronic Series and Periodic Trends

• Isoelectronic Series: Ions/atoms with the same number of electrons.

• Periodic Trends:

  • Electron Affinity: Energy change when an electron is added to an atom.

  • Electronegativity: Tendency of an atom to attract electrons in a bond.

  • Atomic Radii: Size of an atom; decreases across a period, increases down a group.

  • Ionic Radii: Size of an ion; cations are smaller, anions are larger than parent atoms.

  • First Ionization Energy: Energy required to remove the first electron from an atom.

  • Metallic Character: Increases down a group, decreases across a period.

  • Anomalies in Trends: Some elements deviate from general trends due to electron configurations.

  • Acidity: Related to the ability to donate H+; increases across a period and down a group for oxoacids.

Concept	Definition	Key Equation (if applicable)
Molarity (M)	Moles of solute per liter of solution	$M = \frac{\text{mol}}{\text{L}}$
Ideal Gas Law	Relates P, V, n, T for an ideal gas	$PV = nRT$
Percent Yield	Efficiency of a reaction	$\text{Percent Yield} = \frac{\text{Actual Yield}}{\text{Theoretical Yield}} \times 100\%$
Graham’s Law	Relative rates of effusion	$\frac{\text{Rate}_1}{\text{Rate}_2} = \sqrt{\frac{M_2}{M_1}}$
Bohr Equation	Energy change for electron transitions	$\Delta E = B \left( \frac{1}{n_f^2} - \frac{1}{n_i^2} \right)$
Rydberg Equation	Wavelengths of hydrogen emission lines	$\frac{1}{\lambda} = R \left( \frac{1}{n_1^2} - \frac{1}{n_2^2} \right)$

Additional info: This guide expands on the review topics by providing definitions, equations, and examples for each concept, ensuring a comprehensive overview for exam preparation.