Chapter 7: Chemical Reactions and Quantities – Study Notes

Chapter 7: Chemical Reactions and Quantities

Learning Objectives Overview

This chapter focuses on the quantitative aspects of chemical reactions, including balancing equations, identifying reaction types, and performing calculations involving moles, masses, and energy changes. Mastery of these concepts is essential for understanding how chemical reactions occur and how to predict the amounts of substances involved.

Writing and Balancing Chemical Equations

Chemical equations represent the reactants and products in a chemical reaction. Balancing these equations ensures the law of conservation of mass is obeyed.

• Balanced Equation: An equation where the number of atoms of each element is the same on both sides.

• Steps to Balance:

  1. Write the unbalanced equation.

  2. Count atoms of each element on both sides.

  3. Add coefficients to balance atoms.

  4. Check your work.

• Example: $\ce{2H_2 + O_2 \rightarrow 2H_2O}$

Types of Chemical Reactions

Chemical reactions can be classified into several types based on the changes that occur.

• Combination (Synthesis): Two or more substances combine to form one product.

• Decomposition: One substance breaks down into two or more products.

• Single Replacement: One element replaces another in a compound.

• Double Replacement: Two compounds exchange ions to form new compounds.

• Combustion: A substance reacts with oxygen, releasing energy.

• Example: $\ce{CH_4 + 2O_2 \rightarrow CO_2 + 2H_2O}$ (combustion)

Oxidation and Reduction (Redox Reactions)

Redox reactions involve the transfer of electrons between substances.

• Oxidation: Loss of electrons.

• Reduction: Gain of electrons.

• Identifying Redox: Track changes in oxidation numbers.

• Example: $\ce{Zn + Cu^{2+} \rightarrow Zn^{2+} + Cu}$

Avogadro's Number and the Mole

The mole is a counting unit for atoms, molecules, or ions. Avogadro's number defines the number of particles in one mole.

• Avogadro's Number: $6.022 \times 10^{23}$ particles/mole

• Mole: The amount of substance containing Avogadro's number of entities.

• Example: 1 mole of $\ce{H_2O}$ contains $6.022 \times 10^{23}$ molecules.

Molar Mass and Conversions

Molar mass is the mass of one mole of a substance, expressed in grams per mole. It allows conversion between mass and moles.

• Molar Mass Formula: $\text{Molar mass} = \frac{\text{mass (g)}}{\text{moles}}$

• Conversion: Use molar mass to convert grams to moles and vice versa.

• Example: $\ce{H_2O}$ molar mass = 18.02 g/mol

Stoichiometry: Mole Ratios and Calculations

Stoichiometry uses balanced equations to relate quantities of reactants and products.

• Mole Ratio: The ratio of moles of one substance to another in a reaction, from coefficients in the balanced equation.

• Example: In $\ce{2H_2 + O_2 \rightarrow 2H_2O}$, the mole ratio of $\ce{H_2}$ to $\ce{O_2}$ is 2:1.

• Calculations: Use mole ratios to determine amounts of reactants or products.

Limiting Reactant and Theoretical Yield

The limiting reactant is the substance that is completely consumed first, limiting the amount of product formed.

• Limiting Reactant: The reactant that determines the maximum amount of product.

• Theoretical Yield: The maximum amount of product possible, calculated from the limiting reactant.

• Percent Yield Formula: $\text{Percent yield} = \frac{\text{actual yield}}{\text{theoretical yield}} \times 100\%$

• Example: If 10 g of $\ce{A}$ produces 8 g of $\ce{B}$, but the theoretical yield is 10 g, percent yield is 80%.

Energy Changes in Chemical Reactions

Chemical reactions can absorb or release energy, classified as endothermic or exothermic.

• Exothermic Reaction: Releases energy (heat) to the surroundings.

• Endothermic Reaction: Absorbs energy (heat) from the surroundings.

• Identifying: Examine the energy diagram or reaction enthalpy ($\Delta H$).

• Example: Combustion is exothermic; photosynthesis is endothermic.

Summary Table: Key Concepts in Chemical Reactions and Quantities

Concept	Definition	Example
Balanced Equation	Equal number of atoms on both sides	$\ce{2H_2 + O_2 \rightarrow 2H_2O}$
Mole	Counting unit, $6.022 \times 10^{23}$ particles	1 mole $\ce{NaCl}$ = $6.022 \times 10^{23}$ formula units
Molar Mass	Mass of 1 mole (g/mol)	$\ce{CO_2}$ = 44.01 g/mol
Limiting Reactant	Reactant used up first	2 mol $\ce{H_2}$ + 1 mol $\ce{O_2}$: $\ce{O_2}$ is limiting if only 1 mol present
Percent Yield	Actual/Theoretical x 100%	8 g actual / 10 g theoretical = 80%
Exothermic	Releases energy	Combustion
Endothermic	Absorbs energy	Photosynthesis