Chapter 4: Chemical Reactions and Quantities – Study Notes

Chapter 4: Chemical Reactions and Quantities

4.2 Chemical Reactions

Chemical reactions describe the transformation of reactants into products. Each reaction is represented by a chemical equation, which summarizes the substances involved and their changes.

• Reactants: Substances that undergo change during a reaction.

• Products: Substances formed as a result of the reaction.

• Example: $\mathrm{HCO_3^- (aq) + H^+ (aq) \rightarrow CO_2 (g) + H_2O (aq)}$

4.2 Chemical Equations

Chemical equations provide essential information about a reaction, including the formulas and states of reactants and products, and their relative quantities.

• Formulas of reactants and products indicate the chemical species involved.

• States of reactants and products are denoted by (s), (l), (g), and (aq):

Abbreviation	State
(g)	Gas
(l)	Liquid
(s)	Solid
(aq)	Aqueous (water solution)

• Relative numbers of reactant and product molecules are shown by coefficients.

4.2 Combustion of Methane

The combustion of methane is a classic example used to illustrate the law of conservation of mass and the need to balance chemical equations.

• Unbalanced equation: $\mathrm{CH_4 (g) + O_2 (g) \rightarrow CO_2 (g) + H_2O (g)}$

• Problem: Conservation of mass requires equal numbers of each atom on both sides.

• Example: 4 H atoms on the left, 2 H atoms on the right; equation must be balanced by changing coefficients, not subscripts.

• Balanced equation: $\mathrm{CH_4 (g) + 2 O_2 (g) \rightarrow CO_2 (g) + 2 H_2O (g)}$

4.2 Balancing Chemical Equations

Balancing equations ensures the law of conservation of mass is obeyed. Coefficients are adjusted so that the number of atoms of each element is the same on both sides.

• Example: Balance lead(II) nitrate and potassium chloride to form lead(II) chloride and potassium nitrate.

• Example: Find the stoichiometric coefficient for oxygen in the reaction: $\mathrm{NH_3 (g) + O_2 (g) \rightarrow NO (g) + H_2O (g)}$

4.3 Reaction Stoichiometry: Mole-to-Mole

Stoichiometry allows us to relate quantities of reactants and products using the balanced chemical equation.

• Example: Combustion of ethane: $\mathrm{2 C_2H_6 (g) + 7 O_2 (g) \rightarrow 4 CO_2 (g) + 6 H_2O (g)}$

• Use mole ratios to determine how many moles of product are formed from a given amount of reactant.

• Sample problems:

  • How many moles of CO2 are produced from 35.0 moles of ethane?

  • How many moles of O2 are needed to react with 0.244 moles of ethane?

4.3 Reaction Stoichiometry: Mass-to-Mass

Mass-to-mass conversions use molar masses and stoichiometric coefficients to relate the mass of one substance to the mass of another in a reaction.

• General process:

  1. Convert mass of A to moles of A.

  2. Use mole ratio from balanced equation to find moles of B.

  3. Convert moles of B to mass of B.

• Example: Neutralization of stomach acid by magnesium hydroxide:

  • $\mathrm{Mg(OH)_2 (s) + 2 HCl (aq) \rightarrow 2 H_2O (l) + MgCl_2 (aq)}$

  • Calculate mass of HCl neutralized by 3.26 g Mg(OH)2.

• Example: Photosynthesis: $\mathrm{6 CO_2 (g) + 6 H_2O (l) \rightarrow 6 O_2 (g) + C_6H_{12}O_6 (aq)}$

• Example: Aluminum and manganese dioxide: $\mathrm{3 MnO_2 + 4 Al \rightarrow 3 Mn + 2 Al_2O_3}$

4.4 Stoichiometric Relationships

In many reactions, one reactant is completely consumed first, stopping the reaction. This reactant is called the limiting reactant.

• Limiting reactant: The reactant that is used up first and limits the amount of product formed.

• Excess reactant: The reactant that remains after the reaction is complete.

• Example: $\mathrm{A + B \rightarrow C}$

4.4 Calculating Limiting Reactant, Theoretical Yield, and Percent Yield

To determine the limiting reactant, compare the amount of product each reactant could produce. The smallest amount indicates the limiting reactant. Theoretical yield is the maximum amount of product possible, actual yield is what is obtained, and percent yield measures efficiency.

• Example: Combustion of methane:

  • $\mathrm{CH_4 (g) + 2 O_2 (g) \rightarrow CO_2 (g) + 2 H_2O (g)}$

  • 1 mol CH4 reacts with 2 mol O2.

  • If 5 mol CH4 and 8 mol O2 are present, O2 is limiting (enough for 4 mol CO2).

• Definitions:

  • Theoretical yield: Amount of product based on limiting reactant.

  • Actual yield: Amount of product actually produced.

  • Percent yield: $\displaystyle \frac{\text{actual yield}}{\text{theoretical yield}} \times 100\%$

• Sample problems:

  • Given masses of H2 and N2, determine limiting and excess reagent, theoretical yield, and percent yield for ammonia synthesis.

  • Given percent yield and mass of titanium desired, calculate required mass of TiO2 for titanium production.

  • Given masses of Mg and O2, determine limiting reactant and theoretical yield of MgO.

  • Given masses of Na and Cl2, calculate theoretical yield of NaCl.

Summary Table: Key Stoichiometric Concepts

Term	Definition	Formula/Example
Limiting Reactant	Reactant consumed first, limits product	Compare moles of product from each reactant
Theoretical Yield	Maximum product from limiting reactant	Use stoichiometry
Actual Yield	Measured product from experiment	Lab result
Percent Yield	Efficiency of reaction	$\displaystyle \frac{\text{actual yield}}{\text{theoretical yield}} \times 100\%$

Examples of Chemical Reactions

Chemical reactions can be classified and analyzed using the principles above. Practice problems and real-world applications reinforce understanding of stoichiometry and yield calculations.

Additional info: These notes expand on the provided slides by including definitions, stepwise explanations, and sample equations for clarity and completeness.