Thermodynamics and Enthalpy

• Endothermic: $\Delta H > 0$, system absorbs energy
• Exothermic: $\Delta H < 0$, system releases energy
• State function: $\Delta H$ depends only on initial and final states

$q = mc\Delta T$ $q_{rxn} = -q_{solution}$ $\Delta H = \frac{q_{rxn}}{n}$

Coffee cup calorimeter: constant pressure → measures $\Delta H$ Bomb calorimeter: constant volume → measures $\Delta U$ (≈ $\Delta H$ for solutions)

$\Delta H_{rxn} = \sum \Delta H_f^\circ(\text{products}) - \sum \Delta H_f^\circ(\text{reactants})$

Also: can sum individual reaction enthalpies to find the overall $\Delta H$ (Hess cycles).

Standard enthalpy of formation ($\Delta H_f^\circ$): enthalpy to form 1 mol of compound from elements in standard states. Elements in standard state: $\Delta H_f^\circ = 0$.

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

Bond energies are averages → approximate results.

Exothermic: products lower than reactants; $\Delta H < 0$ Endothermic: products higher; $\Delta H > 0$ Activation energy: energy barrier from reactants to transition state

• $q = mc\Delta T$; $\Delta H = -q/n$
• Hess's law: $\Delta H_{rxn} = \sum \Delta H_f^\circ(\text{prod}) - \sum \Delta H_f^\circ(\text{react})$
• Bond energies: broken − formed (approximate)
• Elements in standard state: $\Delta H_f^\circ = 0$