Chemistry · Unit 6 Thermodynamics · 14 min read · Updated 2026-05-11
Hess's Law — AP Chemistry
AP Chemistry · Unit 6 Thermodynamics · 14 min read
1. What is Hess's Law?★★☆☆☆⏱ 3 min
Hess's law (full name: Hess's law of constant heat summation) states that the total enthalpy change for a chemical reaction is independent of the path taken between initial reactants and final products, and depends only on the enthalpy difference between reactants and products. This is a direct consequence of enthalpy being a state function, a core principle of thermodynamics.
The law allows us to calculate ΔH for reactions that cannot be measured directly in a lab, such as reactions with very high activation energy or competing side reactions, by combining ΔH values from other known reactions. On the AP Chemistry exam, Hess's law accounts for ~3-5% of your total score, appearing in both multiple-choice and free-response sections.
2. Manipulating Source Reactions to Calculate ΔH★★★☆☆⏱ 5 min
To solve a basic Hess's law problem, you start with a target reaction (whose ΔH you need to find) and a set of source reactions with known ΔH values. Three valid manipulations are allowed, each with a corresponding change to ΔH:
**Reverse a reaction**: Reversing a reaction flips the direction of energy flow, so the sign of ΔH is flipped.
**Scale stoichiometry**: Multiplying/dividing all coefficients by a constant requires scaling ΔH by the same constant, since enthalpy is an extensive property.
**Add reactions**: After modification, add reactions and cancel common intermediate species that do not appear in the target. Sum the modified ΔH values to get the total ΔH.
Exam tip: When checking for cancellation, always cross off one mole of a species on the left for every one mole on the right. If you end up with a partial mole of an intermediate left over, you likely forgot to scale a source reaction to match the target stoichiometry.
3. ΔH from Standard Enthalpies of Formation★★★☆☆⏱ 4 min
A common AP exam application of Hess's law is calculating the standard enthalpy of reaction from standard enthalpies of formation. By definition, $\Delta H^\circ_f = 0$ for any element in its standard state.
Using Hess's law, any reaction can be broken into two steps: (1) decompose all reactants into their constituent elements (reverse of formation, so ΔH is negative sum of reactant $\Delta H^\circ_f$), (2) combine elements to form products (sum of product $\Delta H^\circ_f$). This gives the shortcut formula:
\Delta H^\circ_{rxn} = \sum n \Delta H^\circ_f (\text{products}) - \sum m \Delta H^\circ_f (\text{reactants})
Exam tip: Always remember the order is products minus reactants, not the reverse. It is easy to mix up the order under test pressure, so write the formula down before plugging in any values.
4. ΔH from Average Bond Enthalpies★★★★☆⏱ 3 min
Another common application of Hess's law is estimating $\Delta H_{rxn}$ from average bond enthalpies. A bond enthalpy is the energy required to break 1 mole of a specific covalent bond in the gaseous state. Breaking bonds is always endothermic (ΔH positive), and forming bonds is always exothermic (ΔH negative).
Using Hess's law, we split the reaction into two steps: break all bonds in reactants to form gaseous atoms, then form all bonds in products from the atoms. This gives the shortcut formula:
\Delta H_{rxn} = \sum (\text{Bond enthalpies of bonds broken}) - \sum (\text{Bond enthalpies of bonds formed})
Exam tip: Always confirm all species are gaseous when using bond enthalpies. If any species is liquid or solid, you must add the enthalpy of phase change to get the correct total ΔH.
Common Pitfalls
Why: Students often adjust stoichiometry correctly but forget reversing a reaction flips energy flow.
Why: The 'products minus reactants' rule is easy to flip under test pressure.
Why: Students remember to change the equation but forget ΔH is an extensive property that depends on the amount of reactant.
Why: Students often count all bonds instead of only those that change, or miscount C-H bonds in hydrocarbons.
Why: Students memorize that oxygen has ΔHf = 0, but forget this only applies to the element in its standard state.
Why: Students rush to get a numerical answer and miss incorrect intermediate cancellation.