Chemical Reactions and Energetics
What to know
Candidates must classify reactions (synthesis, decomposition, single and double replacement, combustion, acid-base neutralization, redox), balance equations, and predict products using solubility and activity rules. Energetics covers enthalpy, calorimetry (q = mcΔT), Hess's law, bond energies, and spontaneity (ΔG = ΔH − TΔS).
Be able to assign oxidation numbers and identify what is oxidized and reduced. Know that in redox, oxidation is loss of electrons and reduction is gain. Recognize that combustion of hydrocarbons yields CO2 and water, and that double replacement reactions often produce a precipitate, gas, or weak electrolyte.
Traps
- Forgetting to balance oxygen last, or leaving polyatomic ions ungrouped.
- Sign errors in calorimetry and enthalpy: exothermic reactions have negative ΔH, and the surroundings gain heat.
- Assuming a reaction is spontaneous because it is exothermic; entropy term matters at high temperature.
- Misidentifying the oxidizing agent (it is the species reduced, not oxidized).
- Ignoring state symbols; a reaction predicted to form an insoluble solid is driven forward.
Worked example
Calorimetry: 50.0 g of water warms by 10.0 °C (c = 4.18 J/g·°C). q = (50.0)(4.18)(10.0) = 2090 J ≈ 2.09 kJ absorbed by water, so the reaction released about 2.09 kJ (negative for the system).
Rule of thumb for Hess's law: if you reverse a step, flip the sign of ΔH; if you multiply a step, multiply ΔH by the same factor.