Define enthalpy change, classify reactions as exothermic or endothermic, and use calorimetry data with energy profile diagrams.

What this dot point is asking

TASC expects you to define enthalpy change, classify and represent reactions on energy profiles, and calculate molar enthalpy changes from calorimetry data.

Exothermic and endothermic reactions

Every reaction breaks bonds (absorbing energy) and makes bonds (releasing energy); the net result sets the energy change.

• In an exothermic reaction, more energy is released forming new bonds than is absorbed breaking old ones, so heat flows out and the temperature rises. $\Delta H$ is negative.
• In an endothermic reaction, more energy is absorbed breaking bonds than is released, so heat flows in and the temperature falls. $\Delta H$ is positive.

Energy profile diagrams

An energy profile plots enthalpy against reaction progress. Reactants and products sit at different enthalpy levels, and the activation energy $E_a$ is the height of the peak above the reactants. For an exothermic reaction the products lie below the reactants ($\Delta H < 0$); for an endothermic reaction the products lie above ($\Delta H > 0$). The vertical gap between reactants and products is $\Delta H$.

Calorimetry

We measure enthalpy changes by capturing the heat in a known mass of water and recording the temperature change:

Standard enthalpy changes

Chemists define standard enthalpy changes so values can be compared fairly. The standard enthalpy of combustion ($\Delta H_c^\circ$) is the heat released when one mole of a substance burns completely in oxygen under standard conditions. The standard enthalpy of formation ($\Delta H_f^\circ$) is the heat change when one mole of a compound forms from its elements in their standard states; elements in their standard state are defined as zero. The standard enthalpy of neutralisation is the heat released when an acid and base react to form one mole of water, and for strong acid and strong base it is almost constant (about $-57\ \text{kJ mol}^{-1}$) because the reaction is essentially $\text{H}^+ + \text{OH}^- \rightarrow \text{H}_2\text{O}$ in every case.

Sources of experimental error in calorimetry

A school calorimetry result is almost always less exothermic than the data-book value, for predictable reasons: heat is lost to the surroundings and to the apparatus rather than all going into the water; combustion may be incomplete, leaving soot and producing less energy; and some fuel evaporates without burning. Recognising these systematic errors, and stating that they make the measured $\Delta H$ less negative, is frequently worth a mark in TASC questions.

In the exam, show $q = mc\Delta T$ with the correct mass and specific heat, divide by moles for a molar value, attach the sign that matches an exothermic or endothermic process, and explain experimental losses when comparing with a data-book value.