Topic 4: Chemical Reactions

Chemical reactions describe how matter rearranges as bonds break and form. This unit teaches you to represent reactions with balanced equations, predict products by reaction type, and use mole ratios to calculate quantities. You will also learn to track electron transfer in redox reactions and to read solutions through titration.

Writing and Balancing Equations

A chemical equation places reactants on the left and products on the right, separated by an arrow. Because atoms are conserved (the law of conservation of mass), every element must appear in equal numbers on both sides. We balance by adjusting coefficients, the whole numbers placed in front of formulas, never by changing subscripts inside a formula. Changing a subscript would change the identity of the substance. For example, combustion of methane balances as CH4 + 2 O2 -> CO2 + 2 H2O. Count atoms: left has 1 C, 4 H, 4 O; right has 1 C, 4 H, 4 O. A reliable strategy is to balance the most complex molecule first, leave free elements (like O2 or H2) for last, and treat polyatomic ions that survive intact as single units. Always reduce coefficients to the smallest whole-number ratio.

Molecular, Ionic, and Net Ionic Equations

For reactions in aqueous solution, the same reaction can be written three ways. The molecular (or formula) equation shows complete neutral compounds, such as AgNO3(aq) + NaCl(aq) -> AgCl(s) + NaNO3(aq). The complete ionic equation shows every strong electrolyte broken into its dissociated ions: Ag+ + NO3- + Na+ + Cl- -> AgCl(s) + Na+ + NO3-. Ions that appear unchanged on both sides are spectator ions; here Na+ and NO3- are spectators. Removing them gives the net ionic equation, Ag+(aq) + Cl-(aq) -> AgCl(s), which captures the actual chemical change. Net ionic equations must balance both atoms and total charge. Remember that solids, liquids, gases, and weak electrolytes are written in molecular form, while only soluble strong electrolytes are split into ions.

Physical versus Chemical Changes

A physical change alters the form or appearance of matter without producing a new substance: melting ice, dissolving sugar, or grinding a solid all leave the chemical identity intact. A chemical change rearranges atoms into new substances with new properties. Clues that a chemical reaction occurred include formation of a precipitate, release of a gas, a color change, a temperature change without external heating, or emission of light. On the particle level, physical changes affect the spacing and motion of particles or the way they interact, while chemical changes break and form chemical bonds. Be careful: dissolving an ionic compound in water is usually considered a physical process of dissociation, even though ions separate, because no bonds within the ions themselves are broken and the solute can be recovered by evaporation.

Stoichiometry and the Mole

Stoichiometry uses the coefficients of a balanced equation as mole ratios to relate amounts of substances. The mole, with Avogadro's number of 6.022 x 10^23 particles, connects the microscopic and macroscopic worlds. The general path is mass to moles (divide by molar mass), moles of known to moles of unknown (multiply by the mole ratio), then moles to the desired unit. Worked example: how many grams of water form when 8.0 g of methane burns? Moles CH4 = 8.0 / 16.0 = 0.50 mol. The ratio of H2O to CH4 is 2:1, so moles H2O = 1.0 mol. Mass H2O = 1.0 x 18.0 = 18 g. For reactions in solution, moles equal molarity times volume in liters, which lets you carry concentration directly into the same mole-ratio framework.

Limiting Reactant and Percent Yield

When reactants are not mixed in exact stoichiometric proportions, one runs out first and stops the reaction; this is the limiting reactant, and it sets the maximum amount of product. The other reactant is in excess. To find the limiting reactant, convert each reactant to moles, divide by its coefficient, and the smallest value identifies the limiter. Worked example: 2 H2 + O2 -> 2 H2O with 4.0 mol H2 and 3.0 mol O2. H2: 4.0 / 2 = 2.0; O2: 3.0 / 1 = 3.0. Hydrogen is limiting, producing 4.0 mol H2O and leaving 1.0 mol O2 unused. The theoretical yield is the amount predicted by the limiting reactant. Percent yield equals actual yield divided by theoretical yield, times 100, and accounts for losses, side reactions, and incomplete reactions.

Introduction to Titration

Titration is a technique for finding an unknown concentration by reacting a sample with a solution of known concentration, called the titrant, delivered from a buret. The reaction proceeds until the equivalence point, where the moles of titrant exactly satisfy the stoichiometry of the reaction with the analyte. In an acid-base titration the equivalence point is often signaled by an indicator that changes color at the endpoint. The core calculation uses moles = molarity x volume. Worked example: 25.0 mL of HCl is neutralized by 30.0 mL of 0.100 M NaOH. Moles NaOH = 0.0300 L x 0.100 M = 0.00300 mol. The 1:1 ratio gives 0.00300 mol HCl. Molarity HCl = 0.00300 mol / 0.0250 L = 0.120 M. Always work from a balanced equation, since the mole ratio may not be 1:1.

Classifying Reactions

Three reaction types appear constantly in AP Chemistry. Precipitation reactions occur when mixing two solutions forms an insoluble ionic solid; predicting them requires solubility rules, such as the fact that most nitrates and group 1 salts are soluble while many sulfides, carbonates, and hydroxides are not. Acid-base reactions, also called neutralizations, transfer a proton (H+) from an acid to a base, typically producing water and a salt; the net ionic equation for a strong acid with a strong base is H+ + OH- -> H2O. Oxidation-reduction (redox) reactions involve transfer of electrons between species. Recognizing the category lets you predict products and write net ionic equations efficiently. Some reactions belong to more than one category, but precipitation and most simple acid-base reactions are not redox because no oxidation numbers change.

Oxidation Numbers and Identifying Redox

Oxidation numbers are bookkeeping charges assigned by rules: free elements are 0, monatomic ions equal their charge, oxygen is usually -2 (peroxides are -1), hydrogen is +1 with nonmetals and -1 with metals, and the sum equals the overall charge of the species. A reaction is redox if any element's oxidation number changes. The species that loses electrons is oxidized (its oxidation number rises) and acts as the reducing agent; the species that gains electrons is reduced (its number falls) and acts as the oxidizing agent. A memory aid is OIL RIG: Oxidation Is Loss, Reduction Is Gain of electrons. Worked example: in Zn + 2 HCl -> ZnCl2 + H2, zinc goes from 0 to +2 (oxidized) and hydrogen goes from +1 to 0 (reduced), so this single-replacement reaction is redox. Oxidation and reduction always occur together.