Study Guide: Solutions, Colligative Properties, and Chemical Kinetics

Ch. 14: Solids and Modern Materials

Section 14.1: Thirsty Solutions: Why You Shouldn't Drink Seawater

This section introduces the concept of solutions and their importance in everyday life, such as the dangers of drinking seawater. Understanding the basic definitions is essential for studying solution chemistry.

• Solution: A homogeneous mixture of two or more substances.

• Solute: The substance present in a lesser amount, dissolved in the solvent.

• Solvent: The substance present in a greater amount, which dissolves the solute.

• Example: In saltwater, salt is the solute and water is the solvent.

Section 14.2: Types of Solutions and Solubility

This section explores the different types of solutions and the factors that affect their solubility, including entropy and intermolecular forces.

• Types of Solutions: Solutions can be gases, liquids, or solids as either the solute or solvent.

• Entropy: A measure of disorder; solutions form because mixing increases entropy.

• Intermolecular Forces: The most important forces in a solution are solvent-solute, solvent-solvent, and solute-solute interactions.

• Solubility: The maximum amount of solute that can dissolve in a given amount of solvent at a specific temperature.

• Ranking Solubility: Solubility depends on the strength of intermolecular forces between solute and solvent particles.

• Example: Polar solutes dissolve best in polar solvents ("like dissolves like").

Section 14.3: Energetics of Solution Formation

This section discusses the energy changes involved in forming solutions, including enthalpy and hydration.

• Enthalpy of Solution Formation: The total energy change when a solution forms, including breaking and forming intermolecular forces.

• Heat of Lattice Energy: The energy required to separate ions in an ionic solid.

• Heat of Hydration: The energy released when ions are surrounded by water molecules.

• Calculating Heat of Solution: $\Delta H_{solution} = \Delta H_{lattice} + \Delta H_{hydration}$

• Example: Dissolving NaCl in water involves breaking ionic bonds (endothermic) and hydrating ions (exothermic).

Section 14.4: Solution Equilibrium and Factors Affecting Solubility

This section covers the equilibrium between dissolved and undissolved solute and the factors that affect solubility.

• Saturated Solution: Contains the maximum amount of dissolved solute at equilibrium.

• Unsaturated Solution: Contains less solute than can be dissolved at equilibrium.

• Supersaturated Solution: Contains more solute than is stable at equilibrium (unstable).

• Temperature: Solubility of solids generally increases with temperature; for gases, solubility decreases with temperature.

• Pressure (Henry's Law): The solubility of a gas in a liquid is proportional to the partial pressure of the gas above the liquid.

• $C = k_P P_{gas}$

• Example: Carbonated beverages are bottled under high pressure to increase CO2 solubility.

Section 14.5: Expressing Solution Concentration

This section explains how to quantify the amount of solute in a solution using various concentration units.

• Mass Percent: $\text{Mass percent} = \frac{\text{mass of solute}}{\text{mass of solution}} \times 100\%$

• Volume Percent, Mole Fraction, Molality, Molarity: Other common units for expressing concentration.

• Example: A 10% NaCl solution contains 10 g NaCl per 100 g solution.

Section 14.6: Colligative Properties: Vapor Pressure Lowering, Freezing Point Depression, Boiling Point Elevation, and Osmotic Pressure

This section covers properties that depend on the number of solute particles, not their identity.

• Vapor Pressure Lowering (Raoult's Law): $P_{solution} = X_{solvent} P^0_{solvent}$

• Freezing Point Depression: $\Delta T_f = -K_f m$

• Boiling Point Elevation: $\Delta T_b = K_b m$

• Osmotic Pressure: $\Pi = MRT$

• Example: Adding salt to ice lowers its freezing point, causing ice to melt at lower temperatures.

Section 14.7: Colligative Properties of Strong Electrolyte Solutions

This section discusses how strong electrolytes affect colligative properties due to ion dissociation.

• van't Hoff Factor (i): The number of particles a compound produces in solution.

• Expected van't Hoff Factor: For NaCl, $i = 2$ (Na+ and Cl-).

• Colligative Property Equations (with i): $\Delta T_f = -i K_f m$, $\Delta T_b = i K_b m$, $\Pi = i M R T$

Ch. 15: Chemical Kinetics

Section 15.2: The Rate of a Chemical Reaction

This section introduces the concept of reaction rate and how it is measured.

• Reaction Rate: The change in concentration of a reactant or product per unit time.

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

• Balanced Equation: The rate must be determined experimentally and is not always predictable from the equation.

Section 15.3: The Rate Law: The Effect of Concentration on Reaction Rate

This section explains how the rate law relates to reactant concentrations and how to determine the order of a reaction.

• Order of Reaction: The exponents in the rate law indicate the order with respect to each reactant.

• Initial Rate Method: Used to determine the order by varying concentrations and measuring initial rates.

• Example: If doubling [A] doubles the rate, the reaction is first order in A.

Section 15.4: The Integrated Rate Law: The Dependence of Concentration on Time

This section covers how reactant concentrations change over time and introduces integrated rate laws.

• Integrated Rate Laws: Mathematical expressions relating concentration and time for different reaction orders.

• First Order: $[A]_t = [A]_0 e^{-kt}$ or $\ln [A]_t = \ln [A]_0 - kt$

• Second Order: $\frac{1}{[A]_t} = \frac{1}{[A]_0} + kt$

• Half-life: The time required for half the reactant to be consumed. For first order: $t_{1/2} = \frac{0.693}{k}$

Section 15.5: The Effect of Temperature on Reaction Rate

This section discusses how temperature affects reaction rates and introduces the Arrhenius equation.

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

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

• Transition State: The high-energy state during a reaction.

• Collision Theory: Molecules must collide with sufficient energy and proper orientation to react.

Section 15.6: Reaction Mechanisms

This section explains how reactions occur in steps and how to analyze mechanisms.

• Elementary Step: A single step in a reaction mechanism.

• Rate-Determining Step: The slowest step, which controls the overall rate.

• Intermediates: Species produced and consumed during the reaction mechanism.

• Example: The decomposition of hydrogen peroxide occurs in two steps, with an intermediate (HO2).

Section 15.7: Catalysts

This section describes how catalysts increase reaction rates by lowering activation energy.

• Catalyst: A substance that increases the rate of a reaction without being consumed.

• Effect: Provides an alternative pathway with lower activation energy.

• Example: Enzymes are biological catalysts that speed up biochemical reactions.