Energetics

Every chemical reaction involves an energy change. Whether energy is released to the surroundings or absorbed from them depends on the balance between the energy needed to break bonds in the reactants and the energy released when new bonds form in the products.

Exothermic and Endothermic Reactions¶

An exothermic reaction releases heat energy to the surroundings. The temperature of the surroundings rises. The products are at a lower energy than the reactants, so the system has released energy.

An endothermic reaction absorbs heat energy from the surroundings. The temperature of the surroundings falls. The products are at a higher energy than the reactants.

Examples of exothermic reactions: combustion, neutralisation (acid + alkali), respiration, the reaction of sodium hydroxide pellets with water, freezing and condensation (physical).

Examples of endothermic reactions: thermal decomposition (e.g. CaCO₃ → CaO + CO₂), photosynthesis, dissolving potassium nitrate in water, melting and evaporation (physical).

ΔH Notation¶

The enthalpy change (ΔH, pronounced "delta H") of a reaction is the heat energy transferred at constant pressure. It is written with the balanced equation:

$\text{CH}_4\text{(g)} + 2\text{O}_2\text{(g)} \rightarrow \text{CO}_2\text{(g)} + 2\text{H}_2\text{O(l)} \quad \Delta H = -890 \text{ kJ mol}^{-1}$

A negative ΔH means heat is released (exothermic). A positive ΔH means heat is absorbed (endothermic).

Bond Energies: Breaking and Forming¶

All chemical reactions involve:

1. Breaking bonds in the reactants — this always requires energy and is endothermic.
2. Forming bonds in the products — this always releases energy and is exothermic.

The overall energy change is the difference between these two:

$\Delta H = \text{energy to break bonds} - \text{energy released forming bonds}$

• If more energy is released in bond formation than is absorbed in bond breaking: ΔH is negative → exothermic.
• If more energy is absorbed in bond breaking than is released in bond formation: ΔH is positive → endothermic.

Energy Profile Diagrams¶

An energy profile diagram (reaction coordinate diagram) shows how the energy of the system changes as reactants are converted into products. The x-axis represents the progress of the reaction and the y-axis represents the potential energy of the system.

Exothermic Profile¶

Key features:

• Reactants start at a higher energy level than products.
• The curve rises to a peak — the transition state (activated complex) — before falling.
• Activation energy ($E_a$): the difference between the reactant energy level and the peak. This is the energy barrier that must be overcome.
• ΔH: the difference in energy between reactants and products. Negative for exothermic reactions (products lower than reactants).

Endothermic Profile¶

Key features:

• Products end at a higher energy level than reactants.
• The curve still rises over an activation energy barrier before reaching the product level.
• ΔH is positive (products higher than reactants).

Effect of a Catalyst¶

A catalyst provides an alternative pathway with a lower activation energy — the peak of the energy profile is lower. The reactant and product energy levels are unchanged, so ΔH is the same. Only the height of the barrier changes.

Calculating Energy Changes¶

The energy transferred in a reaction can be calculated from temperature changes measured in solution using simple calorimetry:

$q = mc\Delta T$

where:

• $q$ = heat energy transferred (J)
• $m$ = mass of solution in grams (for dilute aqueous solutions, assume density = 1 g cm⁻³, so volume in cm³ = mass in g)
• $c$ = specific heat capacity of water = 4.18 J g⁻¹ °C⁻¹
• $\Delta T$ = temperature change in °C

For an exothermic reaction, the temperature rises ($\Delta T$ is positive) and the reaction releases energy, so $q$ is quoted as negative (by convention, energy is released to the surroundings, which is from the reaction).

For an endothermic reaction, the temperature falls ($\Delta T$ is positive but the reaction absorbs energy).

Standard assumptions in CSEC calculations:

• The density of the solution is 1 g cm⁻³.
• The specific heat capacity is that of pure water (4.18 J g⁻¹ °C⁻¹).
• Heat exchange with the surroundings and the calorimeter is negligible.

Heat of Neutralisation¶

Heat of neutralisation is the enthalpy change when an acid and a base react to produce one mole of water. For strong acid–strong alkali reactions, it is approximately −57 kJ mol⁻¹ regardless of which strong acid and strong alkali are used, because the net reaction is always:

$\text{H}^+\text{(aq)} + \text{OH}^-\text{(aq)} \rightarrow \text{H}_2\text{O(l)} \quad \Delta H = -57 \text{ kJ mol}^{-1}$

Heat of Solution¶

Heat of solution is the enthalpy change when one mole of solute dissolves in excess solvent. Dissolving potassium nitrate in water is endothermic (the solution cools); dissolving sodium hydroxide pellets is exothermic (the solution warms significantly).

Thermometric Titration¶

In a thermometric titration, temperature is recorded as acid is added progressively to alkali. Temperature rises as neutralisation proceeds (exothermic), reaches a maximum at the equivalence point, then falls as excess acid cools the mixture. Plotting temperature against volume of acid added produces two straight lines — the intersection identifies the exact equivalence point, without requiring a colour indicator.