Thermochemistry

Notes

Endothermic and exothermic

• Exothermic: releases heat to the surroundings; $\Delta H < 0$; the surroundings get warmer.
• Endothermic: absorbs heat; $\Delta H > 0$; the surroundings get colder.
• Energy diagrams: products lower than reactants means exothermic.
• Thermal equilibrium: heat flows from hot to cold until temperatures match. The heat lost by one object equals the heat gained by the other.

Calorimetry

$q = mc\Delta T$, where $c$ is specific heat (water: 4.18 J/g·°C).

• In a coffee-cup calorimeter, $q_{reaction} = -q_{solution}$.
• $\displaystyle \Delta H_{rxn} = \frac{q_{reaction}}{\text{mol of limiting reactant}}$, usually in kJ/mol.
• Phase changes: temperature stays constant while energy goes into breaking intermolecular attractions. $q = n\Delta H_{fus}$ or $n\Delta H_{vap}$.

Bond enthalpies

$\displaystyle \Delta H \approx \sum(\text{bonds broken}) - \sum(\text{bonds formed})$

• Breaking bonds absorbs energy; forming bonds releases energy.

Enthalpy of formation and Hess's law

• $\displaystyle \Delta H^\circ_{rxn} = \sum \Delta H^\circ_f(\text{products}) - \sum \Delta H^\circ_f(\text{reactants})$. Elements in their standard state have $\Delta H^\circ_f = 0$.
• Hess's law: if you reverse a reaction, change the sign of $\Delta H$. If you multiply a reaction by $n$, multiply $\Delta H$ by $n$. Add steps to get the target reaction.