General Chemistry Exam 2 Study Guide: Chapters 5, 6, and 8

General Chemistry Exam 2 Study Guide

Overview

This study guide outlines the key topics, concepts, and problem-solving skills required for Exam 2 in a General Chemistry course. The exam will cover material from Chapters 5, 6, and 8, focusing on chemical quantities, molecular structure, and the properties of gases. Students are expected to demonstrate proficiency in both conceptual understanding and quantitative problem-solving using dimensional analysis and conversion factors.

Chapter 4 Review: Foundational Skills

Conversion Factors and Dimensional Analysis

• Understand and use unit conversions, density, and other derived units in chemical calculations.

• Apply dimensional analysis to solve multi-step problems involving chemical quantities.

• Key skills:

  • Converting between mass, moles, and number of particles.

  • Using Avogadro's number ($6.022 \times 10^{23}$ particles/mol).

  • Calculating molar mass (g/mol).

  • Determining empirical and molecular formulas from percent composition data.

• Example: Calculate the number of molecules in 5.00 g of glucose ($C_6H_{12}O_6$).

Chapter 5: Chemical Quantities and Formulas

Key Concepts

• Mole Concept: The mole is a counting unit in chemistry, defined as the amount of substance containing as many entities as there are atoms in 12 g of carbon-12.

• Avogadro's Number: $6.022 \times 10^{23}$ entities/mol.

• Molar Mass: The mass of one mole of a substance, expressed in g/mol.

• Percent Composition: The percentage by mass of each element in a compound.

• Empirical and Molecular Formulas:

  • Empirical formula: simplest whole-number ratio of atoms in a compound.

  • Molecular formula: actual number of atoms of each element in a molecule.

• Key Equations:

  • $\text{Moles} = \dfrac{\text{mass (g)}}{\text{molar mass (g/mol)}}$

  • $\text{Number of particles} = \text{moles} \times 6.022 \times 10^{23}$

  • $\text{Percent composition} = \dfrac{\text{mass of element}}{\text{total mass of compound}} \times 100\%$

• Example: Determine the empirical formula of a compound that is 40.0% carbon, 6.7% hydrogen, and 53.3% oxygen by mass.

Chapter 6: Molecular Structure and Bonding

Lewis Structures and Molecular Geometry

• Lewis Structures: Diagrams that show the bonding between atoms and the lone pairs of electrons in a molecule.

• Drawing Lewis Structures:

  • For binary molecules, molecules with one central atom, simple polyatomic ions, and molecules with multiple bonds.

  • Include expanded octets for elements in period 3 and beyond.

• Resonance Structures: Multiple valid Lewis structures for a molecule or ion that differ only in the placement of electrons.

• Formal Charge: Used to determine the most stable Lewis structure.

• VSEPR Theory: Valence Shell Electron Pair Repulsion theory predicts molecular shapes based on electron pair repulsions.

• Polarity: Determined by molecular geometry and bond dipoles.

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

• Bond Angles: Common bond angles include 109.5°, 120°, and 180°.

• Example: Draw the Lewis structure for $CO_2$ and determine if the molecule is polar or non-polar.

Chapter 8: Gases and Their Properties

Gas Laws and Kinetic Molecular Theory

• Physical Properties of Gases: Gases have low density, are highly compressible, and fill their containers.

• Kinetic Molecular Theory: Explains the behavior of gases in terms of particle motion and energy.

• Gas Pressure: The force exerted by gas particles per unit area.

• Common Units: atm, mmHg, torr, Pa.

• Gas Laws:

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

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

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

  • Ideal Gas Law: $PV = nRT$

• Dalton's Law of Partial Pressures: The total pressure of a gas mixture is the sum of the partial pressures of each component.

• Gas Stoichiometry: Use the ideal gas law to relate moles, volume, pressure, and temperature in chemical reactions involving gases.

• Example: A sample of gas occupies 0.31 L at 372 K and 1.25 atm. What is the pressure if the volume is expanded to 0.83 L at 398 K?

Sample Table: Comparison of Gas Laws

Law	Equation	Variables Held Constant	Relationship
Boyle's Law	$P_1V_1 = P_2V_2$	Temperature, moles	Pressure inversely proportional to volume
Charles's Law	$\dfrac{V_1}{T_1} = \dfrac{V_2}{T_2}$	Pressure, moles	Volume directly proportional to temperature
Avogadro's Law	$\dfrac{V_1}{n_1} = \dfrac{V_2}{n_2}$	Pressure, temperature	Volume directly proportional to moles
Ideal Gas Law	$PV = nRT$	None	Relates all four variables

Problem-Solving Strategies

• Show all work and units in calculations.

• Organize solutions by chapter and topic.

• Use dimensional analysis for multi-step problems.

• Refer to key equations and conversion factors as needed.

• Practice with suggested textbook problems for each chapter.

Additional Info

• Bring a scientific calculator (non-programmable) to the exam.

• Reference sheets with key equations will be provided during the exam.

• Extra paper will be available for problem solving.

• Review the Academic Honesty Statement before the exam.