Physical chemistry

Thermochemistry

Thermochemistry quantifies heat exchanged in reactions, using enthalpy, Hess’s law, and bond energies to estimate thermal effects.

At constant pressure, reaction heat equals the system’s enthalpy change. Keep the sign convention tied to the system: an exothermic reaction has $\Delta H<0$ .

Model and quantities

Read the relation together with assumptions about state, experimental conditions, and sign conventions. Keep units consistent and check dimensions before interpreting a result.

\Delta H_{\mathrm{rxn}} = H_{\mathrm{produits}} - H_{\mathrm{réactifs}} = q_p

Definition: Hess’s law

Because enthalpy is a state function, total $\Delta H$ depends only on initial and final states. Thermochemical equations may be reversed or scaled; reversing a reaction changes the sign of $\Delta H$ .

Quantities in the relation are defined for the reaction or system at hand. In particular, distinguish standard-state quantities from actual conditions and do not infer a mechanism from a general expression alone.

Example: Worked example

Burning 1 mol carbon via C + O₂ → CO₂ has $\Delta H=-394$ kJ/mol. A two-step route with effects −110 and −284 kJ/mol gives the same total, −394 kJ/mol.

Solution

Add the steps: $-110+(-284)=-394$ kJ/mol. No further scaling is needed because the equations already represent one mole of carbon.

Concept summary
Concept	Description	Unit / note
Key relation	Use under stated conditions	Check units and sign convention
Measured quantity	Relates a state or process	Compare data with model
Scope	Model specific conditions	Check assumptions first

Standard formation enthalpies $\Delta H_f^\circ$ allow reaction enthalpy to be calculated by stoichiometric sums: products minus reactants. With bond energies, use bonds broken minus bonds formed.

In the worked example, which result follows from the given data?

Which statement is consistent with this lesson?

References

Peter Atkins, Julio de Paula (2014). Thermochemistry