Energy in Chemical Reactions
Energy Is Always Transferred
Every chemical reaction involves an energy transfer. Energy is never created or destroyed — it is transferred between the reacting chemicals (the system) and the surroundings. Whether a reaction heats up or cools down its environment tells you which direction that transfer goes. This idea underpins everything from hand warmers to sports cold packs to the reactions that power living cells.
An exothermic reaction releases energy to the surroundings — the temperature of the surroundings rises. An endothermic reaction absorbs energy from the surroundings — the temperature of the surroundings falls. A useful memory hook: EXO = energy EXits; ENDO = energy ENters.
Exothermic Reactions in Everyday Life
Combustion is the most familiar exothermic reaction — burning methane in a gas boiler, for example, transfers energy to the water in a central heating system. Neutralisation reactions (acid + alkali) are also exothermic; you can feel the flask warm up. Hand warmers that contain iron powder work by an exothermic oxidation reaction: iron slowly reacts with oxygen from the air and releases heat.
Endothermic Reactions in Everyday Life
Thermal decomposition reactions require continuous heating to proceed because they absorb energy from the surroundings. For example, heating copper carbonate causes it to decompose:
$$CuCO_3 \rightarrow CuO + CO_2$$
Sports injury cold packs are a familiar endothermic application: when the pack is activated, ammonium nitrate dissolves in water in an endothermic process, absorbing heat from the surroundings and cooling the injured area.
Activation Energy and Energy Profile Diagrams
Even an exothermic reaction needs an initial energy input to get started. This is called the activation energy — the energy needed to break existing bonds in the reactants so that new bonds can form in the products. Once a few bonds have been broken, the energy released by bond-making can sustain the reaction.
An energy profile diagram shows this graphically. Exam questions call it a reaction profile, so treat the two names as the same thing. The y-axis is energy; the x-axis is the progress of the reaction (sometimes called the reaction coordinate). The curve rises to a peak — the activation energy — before falling to the energy level of the products.
The activation energy arrow must start at the energy level of the reactants and end at the peak of the curve. A very common error is to draw the arrow from the y-axis or from the products — this scores no marks. Check: does your arrow begin on the reactant plateau and rise to the highest point on the curve?
For an exothermic reaction: the products are at a lower energy level than the reactants. The overall energy change (ΔH) is negative — energy has been released. For an endothermic reaction: the products are at a higher energy level than the reactants. The overall energy change is positive — energy has been absorbed.
On an energy profile diagram, the activation energy arrow always starts at the energy level of the reactants and points up to the peak. This is the most common error in exams. Products below reactants = exothermic; products above reactants = endothermic.
Bond Breaking and Bond Making
The energy changes in reactions are explained at the bond level. Breaking chemical bonds always requires energy input — it is an endothermic process. Forming new chemical bonds always releases energy — it is an exothermic process. The overall energy change of a reaction depends on the balance between these two.
If the energy needed to break bonds in the reactants is less than the energy released when new bonds form in the products, the overall reaction is exothermic. If more energy is needed to break bonds than is released in making them, the reaction is endothermic.
Measuring Energy Changes in the Laboratory
A simple way to investigate whether a reaction is exothermic or endothermic is to measure the temperature change of the solution. If you mix two chemicals in a polystyrene cup, the cup acts as a simple calorimeter. Recording the temperature before and after tells you the direction of energy transfer: a temperature rise confirms the reaction is exothermic; a fall confirms endothermic.
In practice, some energy is always lost to the surroundings even with insulation. Using a lid on the cup, taking readings quickly, and using a digital thermometer all reduce this systematic error. Being aware of the limitations of your equipment and method is part of working scientifically — acknowledging sources of error does not undermine a result, but it does show you understand the difference between what you measured and the true value.
bond energy values are given in kJ/mol and represent the energy needed to break one mole of a particular bond. To calculate the overall energy change:
$$\Delta H = \text{Energy to break bonds in reactants} - \text{Energy released making bonds in products}$$
A negative result means exothermic (energy released); a positive result means endothermic (energy absorbed). Always set out the calculation clearly: list each bond broken with its energy, total them; list each bond made with its energy, total them; then subtract.
Worked example: For \(H_2 + Cl_2 \rightarrow 2HCl\). Bond energies: H–H = 436 kJ/mol; Cl–Cl = 242 kJ/mol; H–Cl = 432 kJ/mol.
Energy to break bonds: 436 + 242 = 678 kJ/mol.
Energy released making bonds: 2 × 432 = 864 kJ/mol.
ΔH = 678 − 864 = −186 kJ/mol. Exothermic.