Enthalpy

What is Enthalpy?

Enthalpy (H) is the total amount of heat energy contained in a system. Enthalpy changes(ΔH) tell us how much heat energy is gained or lost from the system during the reaction. The system is the reaction mixture - for example, at the start of a reaction in solution the system contains only reactant particles, with the solvent and calorimeter acting as the surroundings. A reaction that happens in solution is usually classed as a closed system, as energy can be exchanged but matter generally isn't. 

'Standard enthalpies' are actually enthalpy changes. For example, we take the enthalpy of formation of an element in its standard state to be 0. It means ‘to make an element in its standard state from an element in its standard state, ΔH = 0’. Enthalpy changes can be measured directly through calorimetry experiments or deduced from an energy cycle. This is possible because of Hess’s Law. The standard enthalpy changes in your data booklet (or anywhere else) are all empirical values. Some enthalpy changes can be calculated mathematically by, for example, modelling ions as point charges, giving us theoretical values. Empirical enthalpy changes are usually more accurate than their theoretical counterparts, as all models rely on assumptions and most have limitations greater than the error that is usually introduced by a good calorimetry method. 

Calorimetry

Calorimetry experiments allow us to measure enthalpy changes empirically. The graph below charts temperature during a reaction between magnesium ribbon and hydrochloric acid. Remember, enthalpy changes can be calculated from bond energies, ionisation energies and other enthalpies using Hess's Law or a Born-Haber cycle, so when you are doing these experiments you must compare your result to the literature value, even if you have to calculate the literature value yourself from other published data. The difference between your result and the literature value is known as the error in your experiment. If your experimental error is larger than the uncertainties allowed by your instruments, you have systematic errors. The major source of systematic error in a calorimetry experiment is heat loss. One technique to account for the rate of heat loss and obtain a more accurate value for ∆H is to plot temperature against time, starting from several minutes before the reaction until several minutes after. Then extrapolate the cooling curve to predict the maximum temperature your reaction would have reached if it had not already been losing heat as it occurred. The example below was plotted using Excel, though I had to add the trendlines on another app. 

Bond Energies

Bond enthalpies are positive, as it means 'average energy needed to break 1 mole of gaseous bonds'. For the gaseous diatomic elements (H2, N2, O2, F2, Cl2), bond enthalpy is twice the enthalpy of atomisation. Bond energy calculations are simple but they prove to be some of the trickiest questions in exams. This is not because the calculation is difficult, but because students must draw every reactant and product mentioned in the equation, and you must be familiar with these sorts of questions through practice. Otherwise, you will miscount the number of bonds.

Worked examples and Test Yourself for SL: Bond Enthalpy Calculations, then HL: Enthalpy of Combustion and Formation on the next page here.

Standard Enthalpies

These standard enthalpy definitions should be memorised, as exams often ask for them:

Enthalpy of formation, ΔH⦵f : The enthalpy change when 1 mole of a compound is formed from its elements in their standard states, under standard conditions.

Enthalpy of combustion, ΔH⦵c : The enthalpy change when 1 mole of substance is burnt in excess oxygen, under standard conditions. 

Enthalpy of neutralisation: The enthalpy change when 1 mole of water is formed from the reaction of an acid and a base, under standard conditions. 

These standard enthalpy definitions will not be asked for directly but you should understand and practice using them so that you can construct energy cycles and use Hess’s Law effectively:

Enthalpy of hydration, ΔH⦵hyd : The enthalpy change when 1 mole of gaseous ions are surrounded by water molecules to form aqueous ions in an infinitely dilute solution. If the solvent is not water, the term ‘enthalpy of solvation’ should be used instead. 

Lattice enthalpy, ΔH⦵lat : The enthalpy change when gaseous ions are formed from 1 mole of an ionic compound in its standard state. 

Enthalpy of solution, ΔH⦵sol : The enthalpy change when 1 mole of solute (usually an ionic compound) is dissolved to form an infinitely dilute solution. Notice that  ΔH⦵sol = ΔH⦵lat + ΔH⦵hyd as dissolving can be defined as breaking the lattice of the solid crystal, followed by solvation. 

Enthalpy of atomisation, ΔH⦵at : The enthalpy change when 1 mole of gaseous ions is formed from an element in its standard state. (Bond enthalpy = 2 x ΔH⦵at for gaseous diatomic elements)

Bond enthalpy: The enthalpy change when 1 mole of gaseous covalent bonds is broken homolytically, under standard conditions. (It is usually given as a mean bond energy. Different compounds may contain the same bond, but its enthalpy may not be the same in both compounds.) 

1st electron affinity: The enthalpy change when 1 mole of electrons is added to 1 mole of atoms in the gaseous phase. (You may use the symbol ΔH⦵ea but don’t get it confused with activation energy, which is never written as an enthalpy change and instead given the symbol Ea.)

There are some good (free) virtual labs for calorimetry, like the simulation published by Pearson here.

You can read about the factors that affect enthalpy of solution here.

Try my enthalpy changes quiz here.