Enthalpy

Conceptual Plan for Calorimetry

• Calculate heat absorbed by calorimeter: $q<em>{cal} = C</em>{cal} \times \Delta T$
• Find heat of reaction: $q<em>{rxn} = -q</em>{cal}$
• Calculate internal energy change per mole: $\Delta E<em>{rxn} = \frac{q</em>{rxn}}{n}$

Enthalpy

• Enthalpy, $H$, is defined as: $H = E + PV$
• Change in enthalpy, $\Delta H$, represents heat at constant pressure: $\Delta H<em>{reaction} = q</em>{reaction}$
• $\Delta H$ and $\Delta E$ are generally close, differences significant in gas-heavy reactions.

Exothermic vs Endothermic Reactions

• Exothermic: \Delta H < 0 (heat released)
• Endothermic: \Delta H > 0 (heat absorbed)
• Example heat pack reactions: Iron oxidation releases heat (exothermic); NH4NO3 dissolves absorbs heat (endothermic).

Molecular Perspectives

• Exothermic: Temperature increases; energy from reactant bonds to heat.
• Endothermic: Temperature decreases; surroundings lose heat to product formation.

Stoichiometry and Enthalpy

• Heat change (extensive): Larger reactant masses yield larger $\Delta H$ changes.
• Relationships in reactions affect $\Delta H$ based on stoichiometry.

Measuring $\Delta H$ by Calorimetry

• Calorimetry at constant pressure: $q<em>{reaction} = -q</em>{solution}$.
• Calculate using: $q = m \times C_s \times \Delta T$

Hess’s Law

• Overall $\Delta H$ for a reaction is the sum of the $\Delta H$ from each step if reactions can be expressed in steps.

Standard Conditions

• Standard state: pure gas at 1 atm, pure liquid/solid in stable form at specific temp (usually 25 °C), or 1 M for solutions.
• Standard enthalpy change $\Delta H°$: Enthalpy change with all reactants and products in standard states.
• Standard enthalpy of formation $\Delta H<em>f°$: Enthalpy for formation of 1 mole compound from elements in standard states; pure elements have $\Delta H</em>f° = 0$.