General Chemistry Study Guide: Chemical Reactions, Solutions, and Thermochemistry (Chapters 7-9)

Equations to Memorize

Percent Yield

The percent yield is a measure of the efficiency of a chemical reaction, comparing the actual yield to the theoretical yield.

• Percent Yield Formula:

$\%\ \text{Yield} = \frac{\text{Actual yield}}{\text{Theoretical yield}} \times 100\%$

Solution Dilution

When diluting a solution, the amount of solute remains constant, but the volume changes.

• Dilution Equation:

$M_1V_1 = M_2V_2$

Where: $M$ = molarity, $V$ = volume

Heat Capacity and Bomb Calorimeter

Heat capacity is the amount of heat required to change the temperature of a substance by one degree Celsius.

• Heat Capacity Equation:

$q = C \times \Delta T$

Where: $q$ = heat, $C$ = heat capacity, $\Delta T$ = change in temperature

• Bomb Calorimeter Equation:

$q = C_{cal} \times \Delta T$

Where: $C_{cal}$ = calorimeter heat capacity

Enthalpy from Bond Energies

Bond enthalpy calculations estimate the enthalpy change of a reaction using the energies required to break and form chemical bonds.

• Bond Enthalpy Equation:

$\Delta H = \sum \text{(Bond energies of bonds broken)} - \sum \text{(Bond energies of bonds formed)}$

Sign Convention: Breaking bonds is positive (requires energy); forming bonds is negative (releases energy).

Standard Enthalpy of Reaction ($\Delta H^\circ$)

The standard enthalpy of reaction is calculated using standard enthalpies of formation.

• Standard Enthalpy Equation:

$\Delta H^\circ = \sum n \Delta H^\circ_f \text{(products)} - \sum n \Delta H^\circ_f \text{(reactants)}$

Where: $n$ = stoichiometric coefficient, $\Delta H^\circ_f$ = standard enthalpy of formation

Equations You Will Be Given

• $P = w / t$

• $q = m \times C \times \Delta T$

• $\Delta H = q_p$

• $10^3 \text{ J} = 1 \text{ kJ}$

• $E = q + w$

Chapter 7: Chemical Reactions and Chemical Quantities

Overview

This chapter covers the fundamental concepts of chemical reactions, including how to classify, balance, and quantify reactions.

• Chemical and Physical Changes: Chemical changes result in new substances; physical changes do not alter chemical identity.

• Chemical Equations: Represent chemical reactions using reactants and products.

• Balancing Chemical Equations: Ensure the same number of each atom on both sides of the equation.

• Stoichiometry: Use mole ratios from balanced equations to calculate quantities of reactants and products.

• Limiting Reactant: The reactant that is completely consumed first, limiting the amount of product formed.

• Theoretical Yield: Maximum amount of product possible from given reactants.

• Percent Yield: Actual yield divided by theoretical yield, expressed as a percentage.

• Types of Reactions: Includes precipitation, acid-base, and redox reactions.

• Predicting Products: Use reactivity series for metals and solubility rules for ionic compounds.

Example: In a reaction between sodium and water, sodium is more reactive and produces sodium hydroxide and hydrogen gas.

Chapter 8: Introduction to Solutions and Aqueous Reactions

Overview

This chapter introduces the properties of solutions, calculations involving concentration, and the behavior of substances in aqueous environments.

• Molarity Calculations: Molarity ($M$) is defined as moles of solute per liter of solution.

• Dilution Calculations: Use $M_1V_1 = M_2V_2$ to determine new concentrations after dilution.

• Solution Stoichiometry: Relates volumes and concentrations to moles of reactants and products.

• Electrolytes and Nonelectrolytes: Electrolytes conduct electricity in solution; nonelectrolytes do not.

• Solubility Rules: Predict whether an ionic compound will dissolve in water. (Additional info: Only the first two rows of Table 8.1 are required; exceptions for compounds containing Li+, Na+, K+, NH4+, NO3-, or C2H3O2-.)

• Types of Reactions: Precipitation, acid-base, and redox reactions in aqueous solution.

• Net Ionic Equations: Show only the species that participate in the reaction.

• Acids and Bases: Identify strong and weak acids and bases; predict products of neutralization reactions.

• Redox Reactions: Involve transfer of electrons; identify oxidation and reduction agents.

Example: Mixing solutions of silver nitrate and sodium chloride forms a precipitate of silver chloride.

Chapter 9: Thermochemistry

Overview

Thermochemistry studies the energy changes that occur during chemical reactions, including heat, work, and enthalpy.

• Types of Energy: Kinetic, potential, thermal, and chemical energy.

• First Law of Thermodynamics: Energy cannot be created or destroyed; it can only change forms.

• Heat ($q$) and Work ($w$): $q$ is energy transferred due to temperature difference; $w$ is energy transferred by force.

• Internal Energy ($E$): $E = q + w$

• Enthalpy ($H$): $\Delta H = q_p$ at constant pressure.

• Endothermic vs. Exothermic: Endothermic reactions absorb heat; exothermic reactions release heat.

• Calorimetry: Experimental measurement of heat changes using calorimeters.

• Standard Enthalpy of Formation ($\Delta H^\circ_f$): Enthalpy change for forming 1 mole of a compound from its elements in their standard states.

• Bond Enthalpy Calculations: Use bond energies to estimate $\Delta H$ for reactions.

• Hess's Law: The total enthalpy change for a reaction is the sum of the enthalpy changes for individual steps.

Example: The combustion of methane releases energy as heat, which can be measured using a bomb calorimeter.

Solubility Rules Table (Inferred from context)

Main Purpose: To predict whether an ionic compound will dissolve in water.

Ion/Compound	Solubility	Exceptions
Li+, Na+, K+, NH4+	Soluble	None
NO3-, C2H3O2-	Soluble	None
Cl-, Br-, I-	Soluble	Except with Ag+, Pb2+, Hg22+
SO42-	Soluble	Except with Ba2+, Pb2+, Ca2+, Sr2+

Additional info: Only the first two rows are required for your exam, as noted in the original notes.