A-level Chemistry/AQA/Module 2/Energetics

Energy Change
Energetics is all about the energy transfer involved in the making and breaking of bonds. The breaking of a bond involves the injection of energy into the system so that the bond can no longer hold the parts together; the creation of bonds involves the removal of energy from the system so that components are captured by the binding forces. In a typical reaction some bonds are broken and some are created so some energy is injected and some is removed. Whether the total(net) of these amounts to an injection or removal determines whether the reaction is endothermic (involving an net injection of energy) or exothermic (involving a net removal of energy.)

Enthalpy change (ΔH)
The simple idea described above is complicated by the interaction of the system(reaction) being studied and its environment. This interaction amounts to the changes in volume of the system brought about by the reaction. To deal with this complexity we introduce the idea of enthalpy(given the symbol H). Enthalpy is not measured directly: we concern ourselves instead with CHANGES in enthalpy (ΔH) resulting from a reaction.

For a given reaction change in Enthalpy is equal to the heat energy change under conditions of constant pressure.

Exothermic
Such reactions include the combustion of a fuel like methane CH4 (g) + 2O2 (g) → CO2 (g) + 2H2O (l)               = -890kJ mol−1 Also including the oxicdation of carbohydrates such sa glucose C6H12O6
 * negative ΔH

Endothermic
Such reactions include the thermal decomposition of calcium carbonate CaCO3 (s) → CaO (s) + CO2 (g)                                              = +178kJ mol−1
 * positive ΔH


 * Enthalpy change
 * The heat energy transferred under constant pressure.
 * ΔH = ΔH products – ΔH reactants

We often deal with enthalpy change under standard conditions. i.e.
 * Standard conditions
 * 100kPa, 298K


 * Standard enthalpy of formation
 * $$\Delta H^{\ominus}_f$$
 * The enthalpy change when 1 mole of compound in its standard state is formed from its elements in their standard states (under standard conditions).


 * Standard enthalpy of combustion
 * $$\Delta H^{\ominus}_c$$
 * The enthalpy change when 1 mole of substance is completely burned in oxygen under standard conditions.

Calculating energy changes

 * $$q=mc \Delta T$$
 * heat energy = mass of substance x specific heat capacity x temperature change
 * $$\Delta H = \frac{q}{n}$$
 * enthalpy change = heat energy / moles

Hess's Law

 * Hess's Law
 * The overall enthalpy change is independent of the route the chemical reaction takes place.

Bond Enthalpy

 * Energy/enthalpy needed to break/dissociate a bond averaged over different molecules.
 * Always positive because breaking bonds requires energy.
 * $$\Delta H^\ominus_r = \Sigma \Delta H^\ominus_r (bonds\ broken) - \Sigma \Delta H^\ominus_r (bonds\ made)$$