CHM 109 Exam 1 Study Guide: Solutions, Kinetics, and Intermolecular Forces

Chapter 5: Solutions and Electrolytes

Electrolytes vs Nonelectrolytes

Electrolytes are substances that dissociate into ions when dissolved in water, allowing the solution to conduct electricity. Nonelectrolytes do not dissociate into ions and thus do not conduct electricity.

• Electrolytes: Ionic compounds (e.g., NaCl, KBr) and some acids/bases.

• Nonelectrolytes: Molecular compounds (e.g., sugar, ethanol).

• Conductivity: Only solutions with free ions (electrolytes) carry current.

Dissociation/Ionization Equations for Strong Electrolytes

Strong electrolytes dissociate completely in water. The dissociation equation shows the separation into ions.

• Example: $\mathrm{Na_2S(aq) \rightarrow 2\ Na^+(aq) + S^{2-}(aq)}$

Factors Affecting Solubility

Solubility depends on temperature, pressure (for gases), and the nature of solute and solvent.

• Temperature: Solubility of solids usually increases with temperature; gases decrease.

• Pressure: Only affects gases; higher pressure increases solubility.

• Nature: "Like dissolves like"—polar solutes dissolve in polar solvents.

Solubility Rules in Water

Solubility rules help predict whether a compound will dissolve in water.

• Example: AgCl is insoluble; NaCl is soluble.

• Common rules: All nitrates, acetates, and most alkali metal salts are soluble.

Double Displacement Reactions

Double displacement reactions involve the exchange of ions between two compounds, often forming a precipitate.

• States: Indicate (aq) for aqueous, (s) for solid, (l) for liquid, (g) for gas.

Molecular, Complete Ionic, and Net Ionic Equations

These equations represent reactions at different levels:

• Molecular: Shows compounds as intact units.

• Complete Ionic: Shows all strong electrolytes as ions.

• Net Ionic: Shows only the ions and molecules directly involved in the reaction.

Dilution Equation

Used to calculate the concentration after dilution:

• $C_1V_1 = C_2V_2$

Chapter 12 & 14: Intermolecular Forces and Solution Concentrations

Intermolecular Forces (IMF)

IMFs are forces between molecules affecting boiling points, melting points, and solubility.

• Types: Dispersion (London), dipole-dipole, hydrogen bonding, ion-dipole.

• Boiling Point: Stronger IMFs lead to higher boiling points.

Concentration Units

Several units are used to express concentration:

• Molarity (M): $\mathrm{M = \frac{mol\ solute}{L\ solution}}$

• Molality (m): $\mathrm{m = \frac{mol\ solute}{kg\ solvent}}$

• Mole Fraction (\chi): $\mathrm{\chi = \frac{mol\ component}{total\ mols}}$

• Mole Percent: $\mathrm{\% = \chi \times 100}$

• Percent by Mass: $\mathrm{\% = \frac{mass\ solute}{mass\ solution} \times 100}$

• ppm: $\mathrm{ppm = \frac{mass\ solute}{mass\ solution} \times 10^6}$

Concentration Conversions

Convert between units using molar mass, solution volume, and mass as needed.

Unit Prefixes and Conversions

• Giga (G): $10^9$

• Mega (M): $10^6$

• Kilo (k): $10^3$

• Centi (c): $10^{-2}$

• Milli (m): $10^{-3}$

• Micro (\mu): $10^{-6}$

• Nano (n): $10^{-9}$

Henry’s Law

Describes the solubility of gases in liquids:

• $S_{\text{gas}} = k_H \cdot P$

• Trend: Gas solubility increases with pressure, decreases with temperature.

Raoult’s Law

Relates vapor pressure of a solution to the mole fraction of the solvent:

• $P_{\text{solution}} = \chi_{\text{solvent}} \cdot P^\circ_{\text{solvent}}$

• Effect: Adding solute decreases vapor pressure.

Vapor Pressure

• Temperature: Higher temperature increases vapor pressure.

• IMF Strength: Stronger IMFs decrease vapor pressure.

• Boiling Point: Occurs when vapor pressure equals atmospheric pressure.

Freezing Point Depression and Boiling Point Elevation

• $\Delta T_F = -k_F \cdot m \cdot i$

• $\Delta T_B = k_B \cdot m \cdot i$

• Van’t Hoff Factor (i): Number of particles produced by dissociation (e.g., NaCl $i=2$).

Osmotic Pressure

• $\pi = MRT$

• R: Use $0.082057\ \text{L} \cdot \text{atm} / \text{K} \cdot \text{mol}$

Molar Mass Calculations

Use any concentration equation containing moles to solve for molar mass if grams are given.

Chapter 15: Chemical Kinetics

Kinetics and Rate Definitions

• Kinetics: Study of reaction rates.

• Rate: Change in concentration per unit time.

• Average Rate: Over a time interval.

• Instantaneous Rate: At a specific moment.

Rate Expressions

• For reaction $aA + bB \rightarrow cC + dD$:

• $\text{Rate} = -\frac{1}{a} \frac{d[A]}{dt} = -\frac{1}{b} \frac{d[B]}{dt} = \frac{1}{c} \frac{d[C]}{dt} = \frac{1}{d} \frac{d[D]}{dt}$

Factors Affecting Rate

• Concentration, temperature, presence of catalyst, surface area.

Rate Laws

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

• Order in reactants: Exponents $m$, $n$.

• Overall order: Sum of exponents.

• Units of $k$ depend on overall order.

• Example: $\text{Rate} = k[NO][CO]$ is first order in each, second order overall, $k$ units: $\text{M}^{-1}\text{s}^{-1}$.

Method of Initial Rates

• Compare rates with varying initial concentrations to determine order and $k$.

Integrated Rate Laws and Half-Life

• 0th order: $[A]_t = -kt + [A]_0$

• 1st order: $\ln [A]_t = -kt + \ln [A]_0$

• 2nd order: $\frac{1}{[A]_t} = kt + \frac{1}{[A]_0}$

• Half-life:

  • 0th: $t_{1/2} = \frac{[A]_0}{2k}$

  • 1st: $t_{1/2} = \frac{\ln 2}{k}$

  • 2nd: $t_{1/2} = \frac{1}{k[A]_0}$

Graphs of Integrated Rate Laws

• 0th: $[A]$ vs time, negative slope.

• 1st: $\ln [A]$ vs time, negative slope.

• 2nd: $1/[A]$ vs time, positive slope.

Kinetics and Collision Theory

• Reaction rate depends on frequency and energy of collisions.

• Graph interpretation: Endothermic vs exothermic, transition states, number of steps.

Arrhenius Equation

• $k = A e^{-E_a/(RT)}$

• Variables: $k$ (rate constant), $A$ (frequency factor), $E_a$ (activation energy), $R$ (gas constant), $T$ (temperature in K).

Two-Point Arrhenius Equation

• $\ln \left( \frac{k_2}{k_1} \right) = \frac{E_a}{R} \left( \frac{1}{T_1} - \frac{1}{T_2} \right)$

Arrhenius Plot Information

• Slope of $\ln k$ vs $1/T$ plot is $-E_a/R$.

Molecularity and Rate Laws

• Molecularity: Number of reactant particles in an elementary step.

• Rate law for first step: Only include reactants, not intermediates.

Reaction Mechanisms and Catalysts

• Mechanism: Sequence of elementary steps.

• Catalyst: Lowers activation energy, increases rate, not consumed.

• Intermediate: Formed and consumed during reaction.

• Types of catalysts: Homogeneous, heterogeneous.

Exam Preparation and Requirements

Study Sources

• Textbook, ACS study book, end-of-chapter problems, lecture slides, ALEKS.

Equation Sheet

• Provided equations: Gas conversions, rate laws, half-life, Arrhenius equation.

• Constants: $R = 0.0820574\ \text{L} \cdot \text{atm} / \text{K} \cdot \text{mol}$, $R = 8.3145\ \text{J} / \text{K} \cdot \text{mol}$

Exam Logistics

• Bring scientific calculator, writing utensil (no red/pink/orange ink).

• Show all work for partial credit.

• Follow directions for full credit.

• Significant figures must be used correctly.

Additional Info

• Some questions may come from lecture notes not covered in class.

• No practice exams provided; use listed sources for self-study.