Topic H: Chemical Reactions and Energy

Every reaction involves an energy change, usually as heat exchanged with the surroundings. This part defines exothermic and endothermic changes, represents them on energy-level diagrams, explains the energy balance using bond breaking and making, and shows how calorimetry measures the enthalpy change experimentally.

Exothermic and endothermic reactions

Reactions either release or absorb energy. In an exothermic reaction, energy is given out to the surroundings, so the temperature of the surroundings rises; combustion, neutralisation and many displacement reactions are exothermic. In an endothermic reaction, energy is taken in from the surroundings, so the temperature falls; thermal decomposition and dissolving some salts such as ammonium nitrate are endothermic. The energy change comes from the difference between energy needed to break bonds in the reactants and energy released when new bonds form in the products. If more energy is released than absorbed, the reaction is exothermic; if more is absorbed than released, it is endothermic. Recognising the direction of heat flow is the first step in energetics.

Enthalpy change and its sign

The heat change of a reaction at constant pressure is called the enthalpy change, with the symbol deltaH and units of kJ/mol. By convention, an exothermic reaction has a negative deltaH because the system loses energy to the surroundings, while an endothermic reaction has a positive deltaH because the system gains energy. The size of deltaH tells you how much energy is transferred per mole of reaction. Specific enthalpy changes have names: enthalpy of combustion for burning one mole completely in oxygen, enthalpy of neutralisation for forming one mole of water from acid and alkali, and enthalpy of formation for forming one mole of compound from its elements. Always state both the value and the sign.

Energy-level diagrams and activation energy

An energy-level (energy profile) diagram shows the energy of reactants and products and the path between them. In an exothermic reaction the products lie below the reactants, and the drop equals deltaH (negative). In an endothermic reaction the products lie above the reactants, and the rise equals deltaH (positive). In both cases there is an energy hill between reactants and products: the activation energy is the minimum energy that colliding particles must have for the reaction to start, measured from the reactant level to the top of the hill. A catalyst lowers this hill by providing an alternative pathway, speeding the reaction without changing deltaH, which depends only on the start and end levels.

Bond breaking and bond making

The energy change of a reaction can be understood through bonds. Breaking bonds is endothermic because energy must be supplied to pull bonded atoms apart; making bonds is exothermic because energy is released when atoms come together. Bond enthalpy is the average energy needed to break one mole of a particular bond in the gas phase. The overall enthalpy change can be estimated as deltaH = (total energy to break all bonds in the reactants) - (total energy released making all bonds in the products), or simply bonds broken minus bonds made. A negative result means more energy was released than absorbed, so the reaction is exothermic. Because the values are averages, the estimate is approximate.

Bond enthalpy calculations

To estimate an enthalpy change from bond enthalpies, list every bond broken in the reactants and add their bond enthalpies, then list every bond made in the products and add theirs, and subtract. Worked example: for H2 + Cl2 -> 2 HCl, break one H-H bond (436) and one Cl-Cl bond (242), total 678 kJ absorbed; make two H-Cl bonds (2 x 431 = 862) released. So deltaH = 678 - 862 = -184 kJ/mol, an exothermic reaction. A positive answer would indicate an endothermic reaction. Remember that bond enthalpies apply to gaseous molecules and are averages, so the calculated value will differ slightly from precise experimental data, especially when liquids or solids are involved.

Measuring energy change by calorimetry

Calorimetry measures the heat released or absorbed by a reaction from the temperature change of a known mass of water or solution, using q = m c deltaT, where q is the heat in joules, m is the mass of liquid in grams, c is the specific heat capacity (4.18 J/g/K for water), and deltaT is the temperature change. Dividing by the moles of reactant gives the molar enthalpy change, and the sign is set by whether the temperature rose (exothermic) or fell (endothermic). Worked example: 100 g of water rises by 6.0 K, so q = 100 x 4.18 x 6.0 = 2508 J. Heat lost to the surroundings and evaporation make measured values less exothermic than data-book values.