Energetics
Energy Changes in Reactions
Every chemical reaction involves an energy change. Energy is either transferred to the surroundings or absorbed from them. The key principle is that energy is transferred, never created or destroyed. Two categories describe the direction of energy transfer. Understanding these categories — and the bond-level explanation for why energy is transferred — is fundamental to the whole of chemistry. Energy cannot be created or destroyed, only transferred between the system and the surroundings.
Exothermic reactions transfer energy to the surroundings — the temperature of the surroundings increases. Endothermic reactions take in energy from the surroundings — the temperature of the surroundings decreases.
Exothermic reactions include combustion (burning fuels), neutralisation (acid + alkali), and many oxidation reactions. Everyday examples include hand warmers (the crystallisation of a supersaturated sodium acetate solution) and self-heating food cans. Endothermic reactions include thermal decomposition, the reaction of citric acid with sodium hydrogencarbonate, thermal decomposition of calcium carbonate, and photosynthesis. Sports cold packs are a practical example — the endothermic dissolution of ammonium nitrate in water absorbs heat, cooling the injured area.
Measuring Energy Changes (PAG C5)
You can identify whether a reaction is exothermic or endothermic by measuring the temperature change of the surroundings (usually a solution in a polystyrene cup). If the temperature rises, the reaction is exothermic. If the temperature falls, the reaction is endothermic. A polystyrene cup is used as an insulator to reduce heat loss to the surroundings, giving a more accurate temperature reading. You should record the starting temperature, the maximum (or minimum) temperature, and calculate the change. In a well-insulated system, a bigger temperature change indicates a more energetic reaction.
Reaction Profiles
A reaction profile is a diagram that shows the energy levels of the reactants and products, and the energy change during a reaction. On the vertical axis is energy; on the horizontal axis is the progress of the reaction (sometimes labelled “reaction pathway”). The diagram does not show time — it simply shows the energy changes as reactants transform into products.
For an exothermic reaction, the products are at a lower energy level than the reactants — the energy difference has been released to the surroundings. The profile shows a “hump” from reactants rising to a peak (the activation energy barrier) then falling to a lower product level. For an endothermic reaction, the products are at a higher energy level than the reactants — energy has been absorbed. The profile shows a hump rising to the peak then settling at a level higher than the starting point. In both cases, the hump represents the activation energy — the minimum energy needed for the reaction to start. You must label four things on any reaction profile: the energy level of reactants, the energy level of products, the activation energy (from reactants to the peak), and the overall energy change (from reactants to products).
Activation energy is the minimum energy that colliding particles need for a reaction to occur. It is shown on a reaction profile as the difference between the reactant energy level and the top of the energy barrier (the peak).
Bond Energy Calculations (Higher)
During a chemical reaction, bonds in the reactants must be broken (this requires energy — endothermic) and new bonds in the products are made (this releases energy — exothermic). The overall energy change depends on the balance between these two processes. This is a Higher Tier calculation that requires you to use given bond energy values:
Step 1: Add up the total energy needed to break all the bonds in the reactants. Step 2: Add up the total energy released when all the bonds in the products are formed. Step 3: Calculate the overall energy change:
$$\text{energy change} = \text{energy to break bonds} - \text{energy released making bonds}$$
If the answer is positive, the reaction is endothermic (more energy needed to break bonds than is released making them). If the answer is negative, the reaction is exothermic (more energy released making bonds than needed to break them).
For example, for \(H_2 + Cl_2 \rightarrow 2HCl\): breaking bonds = H–H (436 kJ) + Cl–Cl (242 kJ) = 678 kJ. Making bonds = 2 × H–Cl (431 kJ) = 862 kJ. Energy change = 678 − 862 = −184 kJ. Negative, so exothermic. More energy was released making the H–Cl bonds than was required to break the H–H and Cl–Cl bonds, so the surplus energy is transferred to the surroundings as heat. This is why the overall reaction is exothermic. Remember: you must always consider all bonds broken and all bonds made, not just some of them.
Breaking bonds requires energy (endothermic). Making bonds releases energy (exothermic). If more energy is released making bonds than absorbed breaking them, the reaction is exothermic overall.