Thermochemistry & Hess's Law Cheat Sheet

A printable reference covering enthalpy, calorimetry, Hess's Law, standard enthalpy of formation, and bond enthalpy for grades 11-12.

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Thermochemistry studies heat changes during chemical and physical processes. Students need this cheat sheet to connect energy diagrams, reaction equations, calorimetry data, and enthalpy calculations. It is especially useful for AP Chemistry and grade 11-12 units where signs, units, and reaction direction matter. Hess's Law helps calculate reaction enthalpy even when the reaction cannot be measured directly. The most important idea is that enthalpy change, $\Delta H$ , depends only on initial and final states, not on the pathway. Calorimetry uses $q = mc\Delta T$ or $q = C\Delta T$ to measure heat transfer. Standard enthalpy calculations use $\Delta H^{\circ}_{\mathrm{rxn}} = \sum n\Delta H_f^{\circ}(\mathrm{products}) - \sum n\Delta H_f^{\circ}(\mathrm{reactants})$ . Bond enthalpy estimates use bonds broken minus bonds formed.

Key Facts

At constant pressure, heat flow equals enthalpy change: $q_p = \Delta H$ .
For a temperature change in a substance, heat is calculated by $q = mc\Delta T$ , where $\Delta T = T_f - T_i$ .
For a calorimeter with known heat capacity, heat is calculated by $q_{\mathrm{cal}} = C_{\mathrm{cal}}\Delta T$ .
The system and surroundings exchange equal and opposite heat: $q_{\mathrm{system}} = -q_{\mathrm{surroundings}}$ .
For a chemical reaction, molar enthalpy is calculated by $\Delta H_{\mathrm{rxn}} = \frac{q_{\mathrm{rxn}}}{n}$ .
Hess's Law says that if reactions are added, their enthalpy changes are added: $\Delta H_{\mathrm{total}} = \Delta H_1 + \Delta H_2 + \Delta H_3$ .
Standard reaction enthalpy is calculated by $\Delta H^{\circ}_{\mathrm{rxn}} = \sum n\Delta H_f^{\circ}(\mathrm{products}) - \sum n\Delta H_f^{\circ}(\mathrm{reactants})$ .
Bond enthalpy estimates use $\Delta H_{\mathrm{rxn}} \approx \sum D(\mathrm{bonds\ broken}) - \sum D(\mathrm{bonds\ formed})$ .

Vocabulary

Enthalpy: Enthalpy, $H$ , is the heat content of a system at constant pressure.
Enthalpy change: Enthalpy change, $\Delta H$ , is the heat absorbed or released during a process at constant pressure.
Exothermic reaction: An exothermic reaction releases heat to the surroundings and has $\Delta H < 0$ .
Endothermic reaction: An endothermic reaction absorbs heat from the surroundings and has $\Delta H > 0$ .
Hess's Law: Hess's Law states that the total $\Delta H$ for a reaction is the same no matter how many steps are used.
Standard enthalpy of formation: Standard enthalpy of formation, $\Delta H_f^{\circ}$ , is the enthalpy change when $1\ \mathrm{mol}$ of a compound forms from its elements in their standard states.

Common Mistakes to Avoid

Forgetting to change the sign when reversing a reaction, which is wrong because reversing the reaction changes $\Delta H$ to $-\Delta H$ .
Forgetting to multiply $\Delta H$ when multiplying a reaction, which is wrong because enthalpy is extensive and scales with the coefficients.
Using $q = mc\Delta T$ with Celsius and Kelvin mixed incorrectly, which is wrong because $\Delta T$ has the same size in $^{\circ}\mathrm{C}$ and $\mathrm{K}$ but actual temperatures should not be substituted randomly.
Writing the calorimetry sign backward, which is wrong because if the water gains heat then the reaction loses heat, so $q_{\mathrm{rxn}} = -q_{\mathrm{water}}$ .
Subtracting formations in the wrong order, which is wrong because the correct formula is $\Delta H^{\circ}_{\mathrm{rxn}} = \sum n\Delta H_f^{\circ}(\mathrm{products}) - \sum n\Delta H_f^{\circ}(\mathrm{reactants})$ .

Practice Questions

1 A $150.0\ \mathrm{g}$ sample of water warms from $22.0^{\circ}\mathrm{C}$ to $31.5^{\circ}\mathrm{C}$ . Using $c = 4.184\ \mathrm{J\ g^{-1}\ ^{\circ}C^{-1}}$ , calculate $q$ .
2 A reaction releases $2.50\ \mathrm{kJ}$ of heat when $0.0500\ \mathrm{mol}$ of reactant is consumed. Calculate $\Delta H_{\mathrm{rxn}}$ in $\mathrm{kJ\ mol^{-1}}$ .
3 Use $\Delta H_f^{\circ}$ values to find $\Delta H^{\circ}_{\mathrm{rxn}}$ for $\mathrm{CH_4(g) + 2O_2(g) \rightarrow CO_2(g) + 2H_2O(l)}$ if $\Delta H_f^{\circ}(\mathrm{CH_4}) = -74.8\ \mathrm{kJ\ mol^{-1}}$ , $\Delta H_f^{\circ}(\mathrm{CO_2}) = -393.5\ \mathrm{kJ\ mol^{-1}}$ , $\Delta H_f^{\circ}(\mathrm{H_2O}) = -285.8\ \mathrm{kJ\ mol^{-1}}$ , and $\Delta H_f^{\circ}(\mathrm{O_2}) = 0\ \mathrm{kJ\ mol^{-1}}$ .
4 Explain why Hess's Law allows you to calculate $\Delta H$ for a reaction by adding several chemical equations, even if the actual reaction follows a different pathway.