Subarea IV - Chemical Reactions and Stoichiometry
Subarea IV covers reaction types, balancing, the mole concept, stoichiometric calculations, solutions, acid–base neutralization, and redox. The exam rewards clear unit tracking and correct identification of the limiting reagent.
Reactions and balancing
Balance atoms, then charge for redox. Common types: synthesis, decomposition, single replacement, double replacement, combustion, acid–base, and redox. In redox, oxidation numbers track electron transfer: oxidation is an increase in oxidation number (loss of electrons); reduction is a decrease. The oxidizing agent is reduced; the reducing agent is oxidized.
Mole and stoichiometry
One mole contains 6.022 × 10²³ particles. Molar mass (g/mol) converts mass to moles. Use mole ratios from the balanced equation. For solutions, M = mol solute / L solution; dilution: M₁V₁ = M₂V₂. Percent yield = actual / theoretical × 100%.
Limiting reactant
Convert each reactant to moles of product; the smaller amount is the theoretical maximum. The excess reactant remains after reaction.
Worked example
For 2H₂ + O₂ → 2H₂O, start with 4.0 mol H₂ and 1.5 mol O₂. From H₂: 4.0 mol H₂ × (2 mol H₂O / 2 mol H₂) = 4.0 mol H₂O. From O₂: 1.5 mol O₂ × (2 mol H₂O / 1 mol O₂) = 3.0 mol H₂O. O₂ gives less product, so O₂ is limiting and 3.0 mol H₂O forms. Reasoning: compare product amounts from each reactant; do not simply compare starting moles.
Acid–base and redox
Arrhenius acids produce H⁺ in water; bases produce OH⁻. Brønsted–Lowry acids donate protons; bases accept them. Neutralization: H⁺ + OH⁻ → H₂O. Strong acids and bases dissociate completely; weak ones only partly. In redox, balance mass and charge, often in acidic or basic solution.
Traps
Do not use coefficients as mole ratios before balancing. Do not confuse limiting reactant with the reactant having the smaller mass. Do not forget to convert mL to L for molarity. Do not count spectator ions in net ionic equations.