Revise Measuring enthalpy changes for IB Chemistry (R1.1)

Chemical bonds store energy, and a reaction rearranges those bonds, so almost every reaction transfers energy to or from its surroundings. If the reaction gives out energy — the surroundings get hotter — it is exothermic. Combustion, neutralisation and most oxidation reactions are exothermic; you can feel the heat. If instead the reaction takes energy in — the surroundings get colder — it is endothermic. Dissolving ammonium nitrate, or the thermal decomposition of a carbonate, are endothermic; the flask feels cold. The key is that we judge the direction by what happens to the SURROUNDINGS, because that is what we can actually put a thermometer in.

An energy profile diagram plots enthalpy on the vertical axis against the progress of the reaction on the horizontal axis. The reactants start at one enthalpy level and the products end at another; the difference in height between them is ΔH. Between the two there is a hump — the reaction has to climb over an energy barrier before new bonds can form. The height of that barrier above the reactants is the activation energy, Ea. For an EXOTHERMIC reaction the products sit LOWER than the reactants (the line finishes below where it started, ΔH negative); for an ENDOTHERMIC reaction the products sit HIGHER than the reactants (the line finishes above where it started, ΔH positive).

Enthalpy, symbol $H$, is the heat content of a system at constant pressure. We cannot measure the absolute enthalpy of anything, but we can measure the CHANGE that accompanies a reaction, and that is the enthalpy change, $\Delta H$. It is defined as the enthalpy of the products minus the enthalpy of the reactants, which is why the sign falls out automatically: if the products hold less energy, ΔH is negative and the reaction is exothermic.

$\Delta H = H(\text{products}) - H(\text{reactants})$

Because the amount of heat a reaction gives out depends on conditions such as temperature and pressure, chemists quote a STANDARD enthalpy change, written $\Delta H^{\ominus}$, measured under an agreed set of conditions so that values can be compared fairly. ΔH is always quoted per mole of reaction as written in the balanced equation, in kJ mol⁻¹.

Calorimetry measures an enthalpy change indirectly, by measuring the temperature change of the surroundings — usually water. A simple calorimeter is an insulated container, such as nested polystyrene cups with a lid, so that the heat released or absorbed by the reaction goes into (or comes from) the water rather than leaking to the room. We measure how much the water's temperature changes and turn that into a heat quantity, q, using the relationship below.

$q = mc\Delta T$

Here $q$ is the heat transferred (J), $m$ is the mass of the SURROUNDINGS being heated or cooled — the water — in grams, $c$ is the specific heat capacity (4.18 J g⁻¹ K⁻¹ for water, from the data booklet), and $\Delta T = T_{final} - T_{initial}$ is the temperature change in K (a change of 1 °C is a change of 1 K, so either unit works). Two steps then turn q into a molar enthalpy change: first, the heat gained by the water equals the heat lost by the reaction but with the opposite sign, so $\Delta H = -q/n$; second, divide by the number of moles of the reactant that actually reacted and convert from J to kJ.

A reaction is really the breaking of some bonds and the making of others. Breaking a bond always COSTS energy (it is endothermic), and making a bond always RELEASES energy (it is exothermic). The bond enthalpy is the energy needed to break one mole of a given bond in the gaseous state, and the data booklet lists average values. If the bonds made are stronger, in total, than the bonds broken, the reaction releases a net surplus of energy and is exothermic. This gives a quick way to estimate ΔH without any experiment.

$\Delta H = \Sigma(\text{bond enthalpies of bonds broken}) - \Sigma(\text{bond enthalpies of bonds formed})$

This is the showcase for a calculation topic. In Paper 2 the marks are analytic: each mark is tied to a specific line of working (a method mark, M, or an answer mark, A), and — crucially — the method marks and error-carried-forward (ECF) mean a wrong number early on does not cost you every mark that follows. That is only true if the method is written down. Study how each of the four marks below is earned by a specific line.

Enthalpy is the first step into thermodynamics, the study of what drives reactions. From here, Hess's law lets you combine enthalpy changes you can measure to find ones you cannot; enthalpy cycles use exactly the bond-breaking and bond-making bookkeeping you met above. Later you will add entropy and Gibbs free energy to decide not just how much energy a reaction moves but whether it happens at all. The habit built here — measure the heat, divide by the moles, attach the sign — is the foundation for all of it.