Chapter 8: Chemical Reactions – Structured Study Notes

Chemical Reactions

Introduction

Chemical reactions are fundamental processes in chemistry, involving the transformation of substances through the breaking and forming of chemical bonds. This chapter covers the interpretation and writing of chemical equations, balancing equations, patterns of chemical reactivity, combustion analysis, stoichiometric calculations, limiting reactants, and reaction yields.

Writing and Balancing Chemical Equations

Interpreting and Writing Chemical Equations

A chemical equation uses formulas to express the identities and quantities of substances involved in a physical or chemical change. Reactants are written on the left, products on the right, and a yield arrow points from reactants to products. The equation must be balanced so that the same number and type of each atom appear on both sides, obeying the law of conservation of mass.

• Reactants: Substances consumed in the reaction.

• Products: Substances formed in the reaction.

• States of Matter: Indicated by (g) for gas, (l) for liquid, (s) for solid, and (aq) for aqueous.

Balancing Chemical Equations

Balancing chemical equations is essential to ensure the law of conservation of mass is obeyed. This is achieved by writing stoichiometric coefficients to the left of chemical formulas. The process often involves trial-and-error and is facilitated by changing coefficients of compounds before elements, treating polyatomic ions as units, and carefully counting atoms.

• Do not change chemical formulas; only adjust coefficients.

• Check that all atoms balance after each adjustment.

Example: Balancing the combustion of octane:

• Unbalanced: $\mathrm{C_8H_{18} + O_2 \rightarrow CO_2 + H_2O}$

• Balanced: $\mathrm{2C_8H_{18}(l) + 25O_2(g) \rightarrow 16CO_2(g) + 18H_2O(g)}$

Balancing from Molecular Scenes

Visual representations can help deduce formulas and balance equations. For example, a scene with 4 N2O5 molecules converting to 8 NO2 and 2 O2 molecules is balanced as:

• $\mathrm{2N_2O_5(g) \rightarrow 4NO_2(g) + O_2(g)}$

Patterns of Chemical Reactivity

Types of Reactions

Chemical reactions can be classified into several types:

• Combination Reaction: Two or more reactants form one product. Example: $\mathrm{H_2(g) + Br_2(g) \rightarrow 2HBr(g)}$

• Decomposition Reaction: One reactant forms two or more products. Example: $\mathrm{2KClO_3(s) \rightarrow 2KCl(s) + 3O_2(g)}$

• Combustion Reaction: A compound reacts with O2 to produce CO2 and H2O. Example: $\mathrm{2HCO_2H(l) + O_2(g) \rightarrow 2CO_2(g) + 2H_2O(l)}$

Combustion Analysis

Determination of Empirical Formula

Combustion analysis is used to experimentally determine the empirical formula of a compound. The sample is combusted, and the masses of CO2 and H2O produced are measured to calculate the amounts of C and H in the original sample. The remaining mass is attributed to O.

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

• Molecular formula is determined by dividing the molar mass by the empirical formula mass.

Example: For glucose, empirical formula is CH2O, molecular formula is C6H12O6.

Calculations with Balanced Chemical Equations

Stoichiometric Calculations

Balanced equations allow prediction of product amounts from given reactant quantities. The stoichiometric coefficients indicate the ratios in which substances react and are produced.

• Example: $\mathrm{2CO(g) + O_2(g) \rightarrow 2CO_2(g)}$

• 2 moles of CO are equivalent to 2 moles of CO2.

Worked Example: Urea Synthesis

Urea is synthesized from ammonia and carbon dioxide:

• $\mathrm{2NH_3(g) + CO_2(g) \rightarrow (NH_2)_2CO(aq) + H_2O(l)}$

• Use stoichiometric ratios to calculate product and reactant amounts.

Worked Example: Nitrous Oxide Production

Nitrous oxide is produced by heating ammonium nitrate:

• $\mathrm{NH_4NO_3(s) \rightarrow N_2O(g) + 2H_2O(l)}$

• Use molar masses and stoichiometric ratios to convert between grams and moles.

Limiting Reactants and Reaction Yield

Limiting Reactant

The limiting reactant is the reactant used up first in a reaction, determining the maximum amount of product formed. Excess reactants are present in greater quantities than necessary.

• Example: $\mathrm{CO(g) + 2H_2(g) \rightarrow CH_3OH(l)}$

• Calculate required moles for each reactant to identify the limiting reactant.

Worked Example: Alka-Seltzer Reaction

Alka-Seltzer tablets react with water to produce CO2 gas:

• $\mathrm{3NaHCO_3(aq) + H_3C_6H_5O_7(aq) \rightarrow 3CO_2(g) + 3H_2O(l) + Na_3C_6H_5O_7(aq)}$

• Convert reactant masses to moles, determine limiting reactant, calculate excess reactant and product mass.

Reaction Yield

Theoretical yield is the maximum product formed from the limiting reactant. Actual yield is the amount actually obtained. Percent yield is calculated as:

• $\mathrm{\%\ yield = \frac{\text{actual\ yield}}{\text{theoretical\ yield}} \times 100\%}$

Summary Table: Types of Chemical Reactions

Type	General Form	Example
Combination	A + B → AB	H2 + Br2 → 2HBr
Decomposition	AB → A + B	2KClO3 → 2KCl + 3O2
Combustion	Compound + O2 → CO2 + H2O	2HCO2H + O2 → 2CO2 + 2H2O

Chapter Summary: Key Points

• Interpreting and writing chemical equations

• Balancing chemical equations

• Reaction types: combination, decomposition, combustion

• Combustion analysis and empirical formula determination

• Stoichiometric calculations with balanced equations

• Limiting and excess reactants

• Theoretical and actual yield, percent yield