Enthalpy

Conceptual Plan for Calorimetry

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

Enthalpy

  • Enthalpy, H , is defined as: H = E + PV
  • Change in enthalpy, \Delta H , represents heat at constant pressure: \Delta H{reaction} = q{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{reaction} = -q{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 Hf° : Enthalpy for formation of 1 mole compound from elements in standard states; pure elements have \Delta Hf° = 0 .