General Chemistry Exam 2 Study Guide: Thermochemistry, Gas Laws, and Reactivity

Exam 2 Study Guide Overview

This study guide summarizes key concepts from Chapters 4, 5, and 10 of a General Chemistry course, focusing on solution reactions, thermochemistry, and gas laws. It is designed to help students prepare for multiple-choice questions and problem-solving exercises.

What to Bring

• Scientific calculator

• Pencil and eraser

• Periodic Table

• Constants and reference information (as needed for calculations)

Exam Format

• Multiple choice questions covering concepts and problem-solving

Chapter 4: Solution Reactions and Stoichiometry

Types of Reactions

Understanding different types of chemical reactions in aqueous solutions is essential for predicting products and balancing equations.

• Precipitation reactions: Formation of an insoluble product (precipitate) when two solutions are mixed.

• Acid-base reactions: Transfer of protons between reactants; includes strong and weak acids and bases.

• Redox reactions: Electron transfer between species; oxidation and reduction must be identified.

Identifying Reaction Types

• Determine if a reaction is precipitation, acid-base, or redox by examining reactants and products.

• Use solubility rules to predict precipitate formation.

• Detect redox reactions by changes in oxidation numbers.

Balancing Chemical Equations

• Write balanced molecular equations for reactions.

• Include net ionic and complete ionic equations as needed.

Stoichiometry in Solution

• Use molarity ($M$) and volume ($V$) to calculate moles of solute: $n = M \times V$

• Apply stoichiometric relationships to determine quantities of reactants and products.

Titration Problems

• Calculate unknown concentrations using equivalence point data.

• Apply stoichiometry to acid-base titrations.

Chapter 5: Thermochemistry

Key Concepts

Thermochemistry involves the study of energy changes in chemical reactions, including heat, work, and enthalpy.

• First Law of Thermodynamics: Energy cannot be created or destroyed, only transferred. $\Delta E = q + w$

• State functions: Properties that depend only on the state of the system, not the path taken.

• Calorimetry: Measurement of heat changes in chemical reactions.

Specific Heat and Calorimetry

• Calculate heat ($q$) using specific heat ($C$), mass ($m$), and temperature change ($\Delta T$): $q = m \times C \times \Delta T$

• For water, $C = 4.184 \text{ J/g}^\circ\text{C}$

• Heat lost by one substance equals heat gained by another: $q_{\text{system}} = -q_{\text{surroundings}}$

Enthalpy and Hess's Law

• Calculate enthalpy changes ($\Delta H$) for reactions.

• Use Hess's Law to determine $\Delta H$ for overall reactions by combining steps.

• Standard enthalpy of formation problems: $\Delta H_{\text{rxn}} = \sum n \Delta H_f^\circ(\text{products}) - \sum n \Delta H_f^\circ(\text{reactants})$

Energy Calculations

• Potential energy: $PE = mgh$

• Kinetic energy: $KE = \frac{1}{2}mv^2$

Chapter 10: Gas Laws and Kinetic Theory

Gas Laws

Gas laws describe the relationships between pressure, volume, temperature, and amount of gas.

• Ideal Gas Law: $PV = nRT$

• Partial pressures: $P_{\text{total}} = P_1 + P_2 + \ldots$

• Mole fraction: $X_i = \frac{n_i}{n_{\text{total}}}$

• Standard temperature and pressure (STP): $T = 273.15\text{ K}$, $P = 1\text{ atm}$

Gas Stoichiometry

• Relate volumes of gases to moles using molar volume at STP ($1\text{ mol} = 22.4\text{ L}$ at STP).

• Use gas laws to solve for unknowns in chemical reactions involving gases.

Kinetic Molecular Theory

• Average kinetic energy: $KE_{\text{avg}} = \frac{3}{2}RT$

• Root mean square speed: $u_{\text{rms}} = \sqrt{\frac{3RT}{M}}$

• Effusion and diffusion: Rate of effusion is inversely proportional to the square root of molar mass: $\text{rate}_1/\text{rate}_2 = \sqrt{M_2/M_1}$

Reactivity Series of Metals

Classification and Trends

The reactivity series ranks metals by their tendency to react, especially in redox reactions and displacement reactions.

Most Reactive	Less Reactive
Potassium (K)	Hydrogen (H)
Sodium (Na)	Copper (Cu)
Calcium (Ca)	Silver (Ag)
Magnesium (Mg)	Gold (Au)
Aluminum (Al)
Zinc (Zn)
Iron (Fe)
Lead (Pb)

Trend: Chemical reactivity decreases from top to bottom in the series.

Applications

• Predicting displacement reactions

• Understanding corrosion and extraction of metals

Key Equations and Constants

• $PV = nRT$

• $q = mC\Delta T$

• $\Delta E = q + w$

• $PE = mgh$

• $KE = \frac{1}{2}mv^2$

• $u_{\text{rms}} = \sqrt{\frac{3RT}{M}}$

• $\Delta H_{\text{rxn}} = \sum n \Delta H_f^\circ(\text{products}) - \sum n \Delta H_f^\circ(\text{reactants})$

• Gas constant: $R = 0.08206 \text{ L atm/mol K}$ or $R = 8.314 \text{ J/mol K}$

Example Problems

• Calculate the mass of a precipitate formed in a reaction between two solutions.

• Determine the heat absorbed by water when its temperature increases.

• Find the pressure exerted by a gas in a container using the ideal gas law.

• Predict whether a metal will displace another in a solution based on the reactivity series.

Additional info: Some equations and constants were inferred for completeness and clarity.