CHE 1102 Final Exam Review: General Chemistry II Core Concepts

Properties of Gases

Gas Characteristics and Behavior

Gases are one of the fundamental states of matter, characterized by their ability to expand and fill any container, low density, and high compressibility.

• Compressibility: Gases can be compressed much more easily than solids or liquids due to the large distances between particles.

• Expansion: Gases expand to fill the shape and volume of their container.

• Low Density: The density of gases is much lower than that of solids and liquids.

Example: Air in a balloon expands to fill the entire balloon, regardless of its shape.

Gas Laws and Calculations

Partial Pressure in Gas Mixtures

In a mixture of gases, each gas exerts a pressure as if it were alone in the container. This is called partial pressure.

• Dalton's Law of Partial Pressures:

$P_\text{total} = P_1 + P_2 + P_3 + \ldots$

• Finding Partial Pressure: Multiply the mole fraction of the gas by the total pressure.

$P_A = X_A \times P_\text{total}$ where $X_A = \frac{n_A}{n_\text{total}}$

Ideal Gas Law and Related Equations

• Ideal Gas Law:

$PV = nRT$

• Combined Gas Law:

$\frac{P_1 V_1}{T_1} = \frac{P_2 V_2}{T_2}$

• van der Waals Equation: Accounts for non-ideal behavior in real gases.

$\left(P + a\frac{n^2}{V^2}\right)(V - nb) = nRT$

• Standard Temperature and Pressure (STP): 0°C (273.15 K) and 1 atm pressure.

Kinetic-Molecular Theory

This theory explains the behavior of gases based on the motion of their particles.

• Gas particles are in constant, random motion.

• Collisions between particles are elastic (no energy lost).

• The average kinetic energy is proportional to temperature.

Intermolecular Forces (IMFs) and Related Phenomena

Types of Intermolecular Forces

• London Dispersion Forces: Present in all molecules, especially nonpolar ones; weakest IMF.

• Dipole-Dipole Forces: Occur between polar molecules.

• Hydrogen Bonding: Strongest IMF; occurs when H is bonded to N, O, or F.

Example: Water exhibits hydrogen bonding, leading to its high boiling point and meniscus formation.

Vapor Pressure

• Vapor Pressure: The pressure exerted by a vapor in equilibrium with its liquid at a given temperature.

• Higher IMFs result in lower vapor pressure.

Solutions and Solubility

Solubility Concepts

• "Like Dissolves Like": Polar solvents dissolve polar solutes; nonpolar solvents dissolve nonpolar solutes.

• Soluble: Substance dissolves readily in a solvent.

• Insoluble: Substance does not dissolve appreciably.

• Saturated: Maximum amount of solute dissolved.

• Unsaturated: Less than maximum solute dissolved.

• Supersaturated: More than maximum solute dissolved (unstable).

Solution Concentration Calculations

• Molarity (M):

$M = \frac{\text{moles of solute}}{\text{liters of solution}}$

• Percent by Mass:

$\%\,\text{by mass} = \frac{\text{mass of solute}}{\text{mass of solution}} \times 100$

Chemical Kinetics

Rate Laws and Reaction Rates

• Rate Law: Expresses the rate as a function of reactant concentrations.

$\text{Rate} = k[A]^m[B]^n$

• Determine rate law by comparing changes in concentration and rate across trials.

• Order of reaction is the sum of exponents in the rate law.

Factors Affecting Reaction Rate

• Temperature (higher T increases rate)

• Concentration (higher concentration increases rate)

• Catalysts (lower activation energy, increase rate)

Activation Energy

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

• Catalysts lower Ea but do not affect equilibrium position.

Average Rate Calculation

$\text{Average Rate} = -\frac{\Delta [\text{Reactant}]}{\Delta t}$

Reaction Pathway Graphs

• Show energy changes during a reaction.

• Key parts: reactants, products, activation energy peak, intermediates.

Chemical Equilibrium

Equilibrium Concepts

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

• The concentrations of reactants and products remain constant.

Equilibrium Expressions

• For concentrations (Kc):

$K_c = \frac{[C]^c[D]^d}{[A]^a[B]^b}$

• For pressures (Kp):

$K_p = \frac{(P_C)^c (P_D)^d}{(P_A)^a (P_B)^b}$

Reaction Quotient (Q)

• Calculated like K, but with initial concentrations/pressures.

• Compare Q to K to predict direction of shift.

Le Châtelier's Principle

• System at equilibrium responds to disturbances (concentration, temperature, pressure) to restore equilibrium.

• For temperature changes, consider reaction enthalpy (ΔH°).

• For pressure changes, only reactions involving gases are affected.

Catalysts and Equilibrium

• Catalysts speed up both forward and reverse reactions equally; do not change equilibrium position.

Multi-Step Reactions

• Identify intermediates (produced in one step, consumed in another).

• Rate-determining step is the slowest step in the mechanism.

Acids and Bases

Definitions

• Arrhenius: Acids produce H+ in water; bases produce OH-.

• Brønsted-Lowry: Acids donate protons (H+); bases accept protons.

Strong vs. Weak Acids and Bases

• Strong acids/bases: Completely ionize in solution.

• Weak acids/bases: Partially ionize in solution.

Conjugate Acid-Base Pairs

• Acid loses H+ to become its conjugate base.

• Base gains H+ to become its conjugate acid.

Amphiprotic Substances

• Can act as either an acid or a base (e.g., H2O).

pH Calculations

• pH from [H+]:

$pH = -\log_{10}[H^+]$

• pOH from [OH-]:

$pOH = -\log_{10}[OH^-]$

• Relationship: $pH + pOH = 14$ at 25°C

Acid and Base Dissociation Constants

• Ka (acid dissociation constant): Measures strength of an acid.

• Kb (base dissociation constant): Measures strength of a base.

• Relationship:

$K_a \times K_b = 1.0 \times 10^{-14}$

Titration Curves

• Graph of pH vs. volume of titrant added.

• Shape indicates strength of acid and base involved.

Electrochemistry

Redox Reactions

• Oxidation: Loss of electrons.

• Reduction: Gain of electrons.

• Oxidizing Agent: Causes oxidation; is reduced.

• Reducing Agent: Causes reduction; is oxidized.

Electrochemical Cells

Type	Spontaneity	Purpose	Electron Flow
Voltaic (Galvanic) Cell	Spontaneous	Generates electricity from a redox reaction	Anode to Cathode
Electrolytic Cell	Non-spontaneous	Uses electricity to drive a redox reaction	Anode to Cathode

Note: In both cells, oxidation occurs at the anode and reduction at the cathode.

Additional info: This review covers core topics from General Chemistry II, including gases, solutions, kinetics, equilibrium, acids and bases, and electrochemistry. Students should be able to apply these concepts to calculations, identify key properties, and interpret graphs and data.