Energy Changes in Reactions
Energy and Chemical Reactions
Every chemical reaction involves a change in energy. This energy is transferred between the system (the reacting chemicals) and the surroundings. Energy is never created or destroyed during a reaction — it is only transferred. Two categories describe the direction of this transfer, and understanding them is fundamental to chemistry.
Exothermic reactions release energy to the surroundings — the surroundings get warmer. Familiar examples include combustion (burning wood or gas), neutralisation reactions (acid + alkali), and the reactions that power hand warmers. Endothermic reactions absorb energy from the surroundings — the surroundings get colder. Examples include thermal decomposition reactions, and the dissolving of certain salts in water. Sports cold packs are a practical application: when activated, an endothermic reaction absorbs heat from the surroundings (including an injured limb), cooling the area.
Exothermic reactions transfer energy to the surroundings — temperature of surroundings rises. Endothermic reactions take in energy from the surroundings — temperature of surroundings falls. Both involve transfer, never creation or destruction, of energy.
Measuring Energy Changes
You can determine whether a reaction is exothermic or endothermic by measuring the temperature change of the surroundings — typically a solution in a container. If the temperature rises, the reaction is exothermic; if it falls, endothermic. A calorimetry experiment uses a polystyrene cup (a good insulator) to minimise heat loss to the environment, giving a more accurate measurement. You record the starting temperature, the maximum or minimum temperature reached, and calculate the temperature change: ΔT = Tmax − Tstart.
A student adds excess magnesium powder to 50 cm3 of dilute hydrochloric acid in a polystyrene cup. The temperature rises from 19°C to 38°C. The equation for calculating the heat energy released is:
$$Q = mc\Delta T$$
where \(Q\) is heat energy (joules), \(m\) is the mass of the solution (grams), \(c\) is the specific heat capacity of the solution (we assume it is the same as water: 4.18 J/g/°C), and \(\Delta T\) is the temperature change (degrees Celsius).
Step 1 — Mass: 50 cm3 of water has a mass of approximately 50 g. Step 2 — Temperature change: \(\Delta T = 38 - 19 = 19\)°C. Step 3 — Calculate Q: \(Q = 50 \times 4.18 \times 19 = 3971\) J = approximately 3.97 kJ. This is exothermic (temperature rose). The polystyrene cup reduces heat loss, making the result more accurate, but some heat will still escape so the true value is slightly higher.
Reaction Profiles
A reaction profile is a diagram that shows the energy content of the reactants and products, and the energy barrier that must be overcome for the reaction to proceed. The vertical axis represents energy; the horizontal axis represents the progress of the reaction (sometimes called the reaction pathway). The diagram is not a time graph — it shows energy levels, not rates.
Every reaction profile includes a peak — the top of the energy barrier. The difference between the reactant energy level and the peak is the activation energy. This is the energy needed to start breaking bonds in the reactants. For an exothermic reaction, the products are at a lower energy level than the reactants; the overall energy change is negative (energy is released). For an endothermic reaction, the products are at a higher energy level than the reactants; the overall energy change is positive (energy is absorbed).
When drawing or annotating a reaction profile, you must label four things: (1) the reactants energy level (the starting plateau); (2) the products energy level (the ending plateau); (3) the activation energy (the arrow from reactant level up to the peak); (4) the overall energy change (the arrow from reactant level to product level). For exothermic reactions, the product level is lower — the overall energy change arrow points downward. For endothermic reactions, the product level is higher — the arrow points upward.
A common mistake is forgetting that the activation energy always starts from the reactant energy level (not from zero), and that the overall energy change is measured from reactants to products (not from zero to products). Catalysts reduce the activation energy — they provide an alternative reaction pathway with a lower energy barrier — so they lower the peak of the profile without changing the energy levels of reactants or products.
Activation energy is the minimum energy needed for particles to react — shown on a reaction profile as the gap from the reactant level to the peak. Exothermic: products lower than reactants. Endothermic: products higher than reactants. A catalyst lowers the activation energy without changing reactant or product energy levels.
Bond Energies
At the particle level, energy changes arise from the making and breaking of chemical bonds. Every bond has a bond energy. Breaking bonds takes energy in, so breaking bonds is endothermic — you must put energy in to pull the atoms apart. Making bonds releases energy (an exothermic step — energy is given out when new bonds form). The overall energy change of a reaction depends on the balance between these two processes.
Step 1: Count all the bonds broken in the reactants and add up the energy required to break them (using values given in the question). Step 2: Count all the bonds formed in the products and add up the energy released when they form. Step 3: Calculate the overall energy change:
$$\text{energy change} = \text{energy to break bonds} - \text{energy released making bonds}$$
If the result is negative, the reaction is exothermic (more energy released making bonds than required breaking them). If positive, the reaction is endothermic. For example, for \(H_2 + Cl_2 \rightarrow 2HCl\): energy to break bonds = H–H (436 kJ/mol) + Cl–Cl (242 kJ/mol) = 678 kJ/mol. Energy released making bonds = 2 × H–Cl (431 kJ/mol) = 862 kJ/mol. Energy change = 678 − 862 = −184 kJ/mol. Negative, so exothermic.