Experimental measurement of heat changes

• Coffee cup calorimeter: Constant pressure (q = DeltaH)
• Bomb calorimeter: Constant volume (q = DeltaE)

(C):** Energy required to raise temperature by 1 K

• Specific heat capacity (c): Energy for 1 gram by 1 K
• Molar heat capacity: Energy for 1 mole by 1 K

Heat flows spontaneously from hotter to colder object

• Thermal equilibrium: same final temperature
• First Law: $\Delta E = q + w$ (energy conserved)

$q = mc\Delta T$

• q = heat (J)
• m = mass (g)
• c = specific heat (J/g - K)
• DeltaT = temperature change

Enthalpy change ($\Delta H$) for chemical reaction

• Exothermic: releases heat ($\Delta H < 0$)
• Endothermic: absorbs heat ($\Delta H > 0$)

Heat gained by one object equals heat lost by another (insulated system): $q_{calorimeter} + q_{solution} + q_{reaction} = 0$

Phase change: $q = mL$

• L = latent heat (J/g)
• Heat of fusion: solid -> liquid
• Heat of vaporization: liquid -> gas

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

• Breaking bonds: endothermic (+)
• Forming bonds: exothermic (-)

Enthalpy change when 1 mole of compound forms from elements in standard states

• $\Delta H_f degrees$ for elements in standard states = 0

$\Delta H degrees _{rxn} = \sum n\Delta H degrees _f(\text{products}) - \sum n\Delta H degrees _f(\text{reactants})$

Hess's Law: Overall enthalpy change is sum of enthalpy changes for individual steps

• Can combine reactions algebraically
• If reaction reversed, change sign of $\Delta H$
• If reaction multiplied, multiply $\Delta H$ by same factor