1.1.4—Standard enthalpy change (ΔH⦵)

At constant pressure, calculate heat transferred from mass, specific heat capacity, and temperature change, then divide by reacting moles and apply the sign convention. Check standard conditions and units.

Worked calculation: if 100.0 g of solution warms by 6.0 K and $c=4.18\,\mathrm{J\,g^{-1}\,K^{-1}}$, $q(\mathrm{solution})=mc\Delta T=(100.0)(4.18)(6.0)=2.51\times10^3\,\mathrm{J}=+2.51\,\mathrm{kJ}$. Therefore $q(\mathrm{reaction})=-2.51\,\mathrm{kJ}$. If $0.0500\,\mathrm{mol}$ of limiting reactant reacted, $\Delta H=-2.51/0.0500=-50.2\,\mathrm{kJ\,mol^{-1}}$: the negative sign means the reaction released energy. Heat loss or ignored calorimeter heat capacity usually makes the measured magnitude too small.

A standard molar enthalpy change belongs to the balanced reaction as written, with substances in their stated standard states under standard conditions. A classroom calorimetry value is an experimental estimate: report its conditions and uncertainty separately from a data-book or theoretical standard value. Heat loss to the surroundings or ignored calorimeter heat capacity commonly makes the measured magnitude too small.