Combustion is a chemical reaction in which a fuel reacts rapidly with oxygen (O₂), releasing energy as heat and light. In AP Chem, combustion is the go-to example of an exothermic reaction (negative ΔH) and the basis for heat calculations using q = n × ΔH in Topic 6.6.
Combustion is what happens when a fuel reacts with oxygen and dumps energy into the surroundings as heat and light. For hydrocarbons like methane (CH₄) or propane (C₃H₈), complete combustion always produces CO₂ and H₂O. If you can write "fuel + O₂ → CO₂ + H₂O," you can write a combustion equation.
In AP Chemistry, combustion lives in Unit 6 (Thermochemistry) as the textbook exothermic reaction. The enthalpy of combustion is a negative ΔH, which means the system releases heat to the surroundings as reactants convert to products (EK 6.6.A.1 and 6.6.A.2). Here's the part that confuses people at first. The energy doesn't come from "breaking bonds in the fuel." Breaking bonds costs energy. The reaction is exothermic because the bonds formed in CO₂ and H₂O are stronger (lower potential energy) than the bonds broken in the fuel and O₂. The leftover energy escapes as thermal energy.
Combustion is the main vehicle for testing learning objective 6.6.A, which asks you to calculate the heat q absorbed or released by a reaction using moles of reactant and the molar enthalpy of reaction. Nearly every q = n × ΔH problem on the exam is a combustion problem in disguise. You're handed grams of propane or glucose, a balanced combustion equation, and a ΔH in kJ/mol, and you have to convert mass to moles and scale the enthalpy. Combustion also shows up as the source reaction in calorimetry setups (heat from burning a fuel warms a known mass of water) and as context in other units, like fossil fuel combustion producing NO₂ in a 2019 short FRQ. Knowing the combustion pattern cold saves you time everywhere it appears.
Exothermic Reaction (Unit 6)
Combustion is the poster child for exothermic reactions. Its ΔH is always negative, which per EK 6.6.A.1 means heat flows from the system to the surroundings. If a question says "combustion," you can assume heat is released before you do any math.
Enthalpy Change (Unit 6)
The molar enthalpy of combustion (like -2220 kJ/mol for propane) is your conversion factor between moles of fuel burned and heat released. The whole skill in LO 6.6.A is connecting grams → moles → kJ using that one number.
Oxidation (Unit 6 context, redox connection)
Combustion is rapid oxidation. The fuel loses electrons to oxygen, just very fast and with a lot of heat. Rusting iron is the same chemistry running in slow motion, which is a nice way to remember that oxidation doesn't require flames.
Potential Energy and Thermal Energy (Unit 6)
Combustion converts chemical potential energy stored in bonds into thermal energy. Products like CO₂ and H₂O sit at lower potential energy than the reactants, and the difference shows up as heat transferred to the surroundings until thermal equilibrium is reached (EK 6.6.A.2).
Combustion shows up two main ways. First, as a stoichiometry-plus-enthalpy calculation. A typical question gives you 4.50 g of propane and ΔH = -2220 kJ/mol and asks how much heat is released. The move is mass ÷ molar mass to get moles, then multiply by ΔH. Watch the per-mole basis in the balanced equation, because if the equation shows 2 mol of fuel, the listed ΔH may be for 2 mol. Second, as calorimetry. The 2017 short FRQ had a student burn 2-propanol to determine its enthalpy of combustion experimentally, and practice questions follow the same pattern of using combustion heat to warm a measured mass of water (q = mcΔT on the water side, set equal to the heat from the fuel). Conceptual MCQs also ask why combustion releases so much energy. The credited answer is about stronger bonds in the products, not "energy stored in bonds being broken."
All combustion reactions are exothermic, but not all exothermic reactions are combustion. Combustion specifically requires a fuel reacting with O₂ (usually making CO₂ and H₂O). Dissolving NaOH in water or many acid-base neutralizations are exothermic with zero burning involved. On the exam, "combustion" tells you the reaction type and the products; "exothermic" only tells you the sign of ΔH.
Combustion is the rapid reaction of a fuel with O₂ that releases heat and light, and complete hydrocarbon combustion always produces CO₂ and H₂O.
Combustion is always exothermic, so its enthalpy of reaction (ΔH) is always negative and heat flows from the system to the surroundings.
To find heat released, convert the mass of fuel to moles, then multiply by the molar enthalpy of combustion (q = n × ΔH), checking the mole basis of the balanced equation.
Combustion releases energy because the bonds formed in CO₂ and H₂O are stronger than the bonds broken in the fuel and O₂, not because breaking bonds releases energy.
In calorimetry FRQs, the heat from burning a fuel equals the heat absorbed by the water, so you set n × ΔH equal to mcΔT and solve for the unknown.
Combustion is a reaction where a fuel reacts rapidly with oxygen, releasing heat and light. For hydrocarbons like CH₄ or C₃H₈, complete combustion produces CO₂ and H₂O, and the reaction always has a negative ΔH.
Yes. By definition combustion releases heat, so ΔH of combustion is always negative. That's why enthalpies of combustion like -2220 kJ/mol for propane are reported as negative values.
No, and this is a classic AP trap. Breaking bonds always requires energy. Combustion is exothermic because forming the bonds in CO₂ and H₂O releases more energy than breaking the bonds in the fuel and O₂ costs.
Combustion is a type of oxidation, just fast and heat-releasing. Slow oxidation, like iron rusting, involves the same electron transfer to oxygen but without flames or rapid heat release.
Convert grams of fuel to moles, then multiply by the molar enthalpy of combustion. For example, 4.50 g of propane (44.1 g/mol) is 0.102 mol, and 0.102 mol × 2220 kJ/mol gives about 227 kJ released.