General Chemistry Exam Study Guide

Introduction

This study guide summarizes the key topics and concepts for an upcoming General Chemistry exam, focusing on gases, thermodynamics, and calorimetry. It outlines essential definitions, equations, and problem-solving strategies, as well as exam logistics and allowed materials.

Ch. 5: Gases

5.9 Oxidation-Reduction Reactions

• Oxidation-Reduction (Redox) Definition: A chemical reaction involving the transfer of electrons between species.

• Oxidation: Loss of electrons by a species.

• Reduction: Gain of electrons by a species.

• Oxidizing Agent: The species that is reduced (gains electrons).

• Reducing Agent: The species that is oxidized (loses electrons).

• Assigning Oxidation States: Be able to determine oxidation numbers for elements in compounds and ions.

• Identifying Redox Reactions: Recognize if a reaction is a redox process by changes in oxidation states.

• Note: Balancing redox reactions and activity series are not required for this exam.

5.6 Pressure: The Results of Molecular Collisions

• Definition of Pressure: Force exerted per unit area by gas molecules colliding with container walls.

• Barometer: Device used to measure atmospheric pressure.

• Manometer: Device used to measure the pressure of a gas in a container.

• Pressure Units: Common units include atmospheres (atm), pascals (Pa), torr, and mmHg.

• Pressure Conversions: Be able to convert between different pressure units.

5.7 Gas Laws: Boyle's, Charles', Avogadro's Laws

• Boyle's Law: At constant temperature, the volume of a gas is inversely proportional to its pressure. Equation:

• Charles' Law: At constant pressure, the volume of a gas is directly proportional to its absolute temperature. Equation:

• Avogadro's Law: At constant temperature and pressure, the volume of a gas is directly proportional to the number of moles. Equation:

• Ideal Gas Law: Relates pressure, volume, temperature, and moles of a gas. Equation:

• Standard Temperature and Pressure (STP): 0°C (273.15 K) and 1 atm; 1 mol of an ideal gas occupies 22.4 L at STP.

5.8 Applications of the Ideal Gas Law

• Calculate molar volume, density, and molar mass of a gas using the ideal gas law.

• Use the equation to solve for unknown variables.

5.6 Mixtures of Gases and Partial Pressures

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

• Calculate mole fractions and partial pressures in gas mixtures.

5.7 Gases in Chemical Reactions: Stoichiometry Revisited

• Be able to perform stoichiometric calculations involving gases, using the ideal gas law to relate moles and volumes.

5.8 Kinetic Molecular Theory: A Model for Gases

• Understand the postulates of the kinetic molecular theory (KMT) and how they explain the behavior of ideal gases.

• Relate temperature to the average kinetic energy of gas particles. Equation:

5.9 Graham's Law of Effusion and Diffusion

• Effusion: The process by which gas molecules escape through a small hole.

• Diffusion: The mixing of gases due to molecular motion.

• Graham's Law: The rate of effusion of a gas is inversely proportional to the square root of its molar mass. Equation:

• Use Graham's Law to calculate unknown molar masses.

5.10 Real Gases: The Effects of Size and Intermolecular Forces

• Know the van der Waals equation and how it corrects for non-ideal gas behavior: Equation:

• a: Corrects for intermolecular attractions.

• b: Corrects for finite molecular volume.

Ch. 7: Thermochemistry

7.1 Chemical Handwarmers

• Introduction to thermodynamics and thermochemistry.

7.2 The Nature of Energy: Key Definitions

• Energy: The capacity to do work or produce heat.

• Kinetic Energy: Energy due to motion.

• Potential Energy: Energy due to position or composition.

• Thermal Energy: Energy associated with temperature.

• Law of Conservation of Energy: Energy cannot be created or destroyed, only transformed.

7.3 The First Law of Thermodynamics

• First Law: (change in internal energy equals heat plus work).

• Sign Conventions: Heat absorbed by the system (), work done on the system ().

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

• Internal Energy:

7.4 Quantifying Heat and Work

• Heat Capacity (C): Amount of heat required to raise the temperature of an object by 1°C.

• Specific Heat Capacity (c): Amount of heat required to raise 1 gram of a substance by 1°C.

• Heat Equation:

• Work Equation:

• Be able to solve problems involving heat transfer and work during expansion or compression.

7.5 Measuring Delta E for Chemical Reactions: Constant Volume Calorimetry

• Use of bomb calorimeters to measure changes in internal energy at constant volume.

• Calorimetry equation:

7.6 Enthalpy: The Heat Evolved in a Chemical Reaction at Constant Pressure

• Enthalpy (H):

• Change in Enthalpy:

• Exothermic Reaction: Releases heat ().

• Endothermic Reaction: Absorbs heat ().

• Be able to calculate for reactions and interpret sign conventions.

7.7 Constant-Pressure Calorimetry: Measuring Delta Hrxn

• Use of coffee-cup calorimeters to measure enthalpy changes at constant pressure.

• Calorimetry equation:

• Relationship:

7.8 Relationships Involving Delta Hrxn

• Apply Hess's Law to calculate enthalpy changes for reactions by combining known equations.

Exam Logistics and Allowed Materials

• Exam consists of 25 multiple choice questions (+10 calculation-based questions).

• Duration: 50 minutes.

• Bring a non-programmable, non-graphing calculator and a No. 2 pencil.

• You will be provided with a periodic table, equation sheet, and scratch paper.

• No drinks, food, or other items allowed except pencil and calculator.

Table: Gas Laws and Their Relationships