Chemical Reactions
Reactions, stoichiometry, and redox
Write balanced equations, then convert to net ionic form by removing spectator ions. Precipitation depends on solubility rules; a reaction occurs only if an insoluble solid, a gas, or a weak electrolyte or water forms. Acid–base neutralization produces water and often a salt; redox reactions transfer electrons and are balanced by oxidation-number changes or half-reactions.
Oxidation numbers: elements are 0; monatomic ions equal their charge; oxygen is usually −2 (peroxide −1); hydrogen is usually +1 (hydride −1); fluorine is always −1. The sum of oxidation numbers equals the overall charge. Oxidation is loss of electrons (increase in oxidation number), reduction is gain (decrease). A disproportionation reaction has one element both oxidized and reduced.
Stoichiometry traps are common: identify the limiting reactant before computing product mass; molarity times liters gives moles, not grams; percent yield = actual/theoretical × 100. In titration problems, use the balanced mole ratio, not a 1:1 assumption. Combustion of a hydrocarbon with limited oxygen can produce CO rather than CO2.
Worked example: 25.0 mL of 0.100 M AgNO3 reacts with excess NaCl. Moles Ag+ = 0.0250 L × 0.100 mol/L = 2.50 × 10−3 mol. Since Ag+ + Cl− → AgCl(s) is 1:1, mass AgCl = 2.50 × 10−3 mol × 143.32 g/mol = 0.358 g.
Rule of thumb: balance atoms first, then charge; in redox, balance O with H2O and H with H+ in acidic solution.