3.2.1 Enthalpy changes

If the reacting system releases heat to the surroundings, you call the reaction exothermic and ΔH is negative. If the system takes in heat from the surroundings, it is endothermic and ΔH is positive.

I find it useful to keep “system” and “surroundings” separate: an exothermic reaction can make the surroundings warmer because energy has left the reacting system.

• exothermic: energy released, ΔH < 0
• endothermic: energy absorbed, ΔH > 0

When you draw an enthalpy profile, put reactants and products at their correct relative enthalpies. For an exothermic reaction the products finish lower than the reactants; for an endothermic reaction they finish higher.

Draw ΔH between the reactant and product levels, not from the top of the hump. The hump belongs to activation energy.

Activation energy, Eₐ, is the minimum energy needed for a reaction to occur when particles collide. On a profile diagram, you measure it from the reactant enthalpy level up to the top of the energy barrier.

If you add a catalyst to the diagram, draw a lower peak for the catalysed route. Keep the reactant and product enthalpies where they were, so ΔH stays the same.

A standard enthalpy value refers to substances in their standard states under standard conditions. For OCR, use 100 kPa and a stated temperature, usually 298 K unless the question tells you otherwise.

The standard state is the physical state of the substance under those conditions. So at 298 K, water’s standard state is liquid and oxygen’s is O₂(g). State symbols are doing real work in thermochemistry.

• standard pressure: 100 kPa
• temperature stated, commonly 298 K
• each substance in its standard state

There are three definitions here that I’d learn accurately. Standard enthalpy of formation is for forming one mole of a compound from its elements in their standard states. Standard enthalpy of combustion is for one mole of a substance burning completely in oxygen. Standard enthalpy of neutralisation is for an acid and an alkali forming one mole of water.

For a reaction enthalpy, read the equation exactly as written. If you double every coefficient, the enthalpy change doubles; if you reverse the equation, the sign changes.

• ΔfH°: form 1 mol of compound from its elements in standard states
• ΔcH°: burn 1 mol completely in O₂ under standard conditions
• ΔneutH°: acid + alkali form 1 mol H₂O under standard conditions

Start with the temperature change of the surroundings you measured, then calculate the heat transferred using q = mcΔT. In a simple solution calorimetry question, c is usually the specific heat capacity of the solution and m is the mass being warmed or cooled.

Then convert q to a molar enthalpy change for the reacting amount. Keep the sign straight: if the solution warms up, the surroundings gained heat, so the reaction itself released it and ΔH is negative.

A bond enthalpy value tells you the energy needed to break a particular type of covalent bond in gaseous molecules. The data table gives an average because, for example, a C–H bond does not have exactly the same strength in every molecule.

So a bond-enthalpy calculation is usually an estimate. If you compare it with an experimental ΔH, a small difference is not automatically a mistake; the averaged bond values are one reason the answers differ.

For bond enthalpies, I use one reliable sign rule: breaking bonds costs energy; making bonds releases energy. So add the energies of all bonds broken, then subtract the energies released when the product bonds form.

Count the actual bonds in the balanced equation, including any coefficient in front of a molecule. This is where a neat displayed structure can save you from missing a bond.

A Hess cycle gives you two different routes between the same starting and finishing states. Put the target reaction across the top, then choose a common lower or upper level that matches the data you have.

With combustion data, both sides usually point down to the same combustion products. With formation data, both sides can be built from the same elements. I always draw the arrows before doing any maths; the directions tell you which values need their signs reversed.

Once the cycle is drawn, follow one complete route from the same start to the same finish and make its enthalpy change equal to the other route. Reverse an arrow and you reverse the sign; multiply an equation and you multiply its ΔH.

For an unfamiliar cycle, do not hunt for a memorised formula. Use the arrows. If you can account for every step and arrive at the same final state, Hess’s law does the rest.

If the reaction can be carried out cleanly and you can measure a useful temperature change, you can estimate ΔH directly by calorimetry. Measure masses or volumes, record the temperature change, calculate q, then relate that heat to the reacting amount.

Some enthalpy changes are hard or impossible to measure directly. In those cases, use Hess cycles with enthalpy changes that can be measured, such as combustion or formation data. In a practical answer, also think about heat loss, incomplete reaction and whether the apparatus absorbs some of the energy.