Energy Changes and Reaction Profiles
Energy in Chemical Reactions
All chemical reactions involve energy changes. When chemical bonds are broken, energy must be absorbed; when new bonds are formed, energy is released. The overall energy change depends on the balance between these two processes. If more energy is released than absorbed, the reaction is exothermic. If more energy is absorbed than released, it is endothermic.
Bond breaking is always endothermic (energy absorbed); bond making is always exothermic (energy released). The overall reaction type depends on which process releases or absorbs more energy.
Exothermic Reactions
In an exothermic reaction, the energy released when new bonds form is greater than the energy needed to break the original bonds. This surplus energy is transferred to the surroundings as heat, so the temperature of the surroundings rises. Everyday examples include combustion (burning fuels), neutralisation reactions between acids and alkalis, and the oxidation of metals. Hand warmers contain iron powder that reacts exothermically with oxygen from the air, steadily releasing warmth.
When methane burns in oxygen, new C—O and H—O bonds form in the products carbon dioxide and water. These new bonds release more energy than is needed to break the original C—H and O=O bonds. The equation is:
$$CH_4 + 2O_2 \rightarrow CO_2 + 2H_2O$$
The reaction is strongly exothermic — this is why burning fuels release heat and light.
Endothermic Reactions
In an endothermic reaction, the energy needed to break bonds in the reactants is greater than the energy released when new bonds form. Energy is therefore absorbed from the surroundings and the temperature falls. Thermal decomposition reactions are typically endothermic — for example, heating copper carbonate:
$$CuCO_3 \rightarrow CuO + CO_2$$
Sports injury cold packs exploit this principle: ammonium nitrate dissolves endothermically in water, absorbing heat from the injured area and producing an immediate cooling effect.
Exothermic: temperature of surroundings rises (e.g. combustion, neutralisation, hand warmers). Endothermic: temperature of surroundings falls (e.g. thermal decomposition, dissolving ammonium nitrate, cold packs).
Reaction Profiles
A reaction profile (also called an energy profile diagram) shows the energy of the system as reactants are converted to products. The vertical axis shows energy; the horizontal axis shows the ‘progress of reaction’ (sometimes called the reaction coordinate). Every reaction profile has a ‘hump’ — reactants must reach the top of this energy barrier before products can form.
Exothermic profile: The reactants start at a higher energy level than the products. The difference in height between reactants and products is the energy released to surroundings (−ΔH). The ‘hump’ above the reactants represents the activation energy required to start the reaction.
Endothermic profile: The products sit at a higher energy level than the reactants. Energy has been absorbed from the surroundings (+ΔH). The activation energy is the height from the reactants up to the peak of the hump.
A catalyst lowers the height of the activation energy hump, allowing more collisions to be successful at the same temperature — this is shown on a profile as a lower peak, with the reactant and product energy levels unchanged.
Activation Energy
The activation energy is the minimum energy that particles must possess for a collision to result in a reaction. Even strongly exothermic reactions require an initial energy input to break the first bonds and get the reaction started — a match is needed to ignite methane gas, for example. On a reaction profile, activation energy is the vertical distance from the reactant energy level to the peak of the curve.
A catalyst provides an alternative reaction pathway with a lower activation energy. On a reaction profile, this appears as a lower peak. The overall energy change (ΔH) between reactants and products remains identical; only the barrier height changes. This means more collisions at a given temperature have enough energy to react, so the rate increases.
Bond Energy Calculations
Every chemical bond has a characteristic bond energy (in kJ/mol). By summing the bond energies of all bonds broken in the reactants and all bonds made in the products, the overall energy change for a reaction can be calculated:
$$\Delta H = \text{Energy to break bonds} - \text{Energy released making bonds}$$
A positive ΔH means energy is absorbed overall (endothermic); a negative ΔH means energy is released overall (exothermic).
Calculate the energy change for: \(H_2 + Cl_2 \rightarrow 2HCl\)
Bond energies: H—H = 436 kJ/mol; Cl—Cl = 242 kJ/mol; H—Cl = 431 kJ/mol
Step 1 — bonds broken: 1 × H—H (436) + 1 × Cl—Cl (242) = 678 kJ/mol
Step 2 — bonds made: 2 × H—Cl (2 × 431) = 862 kJ/mol
Step 3 — overall change: \(\Delta H = 678 - 862 = -184\) kJ/mol
The negative value confirms the reaction is exothermic — 184 kJ/mol is released to the surroundings.
ΔH = energy to break bonds − energy released making bonds. Negative ΔH = exothermic (energy released); positive ΔH = endothermic (energy absorbed). Always include the sign and units (kJ/mol) in your answer.