Energetics

Objective		Exothermic and endothermic reactions (2h)
1,2	Deduce from data whether heat is released or absorbed in a reaction.
		Investigate/demonstrate temperature changes in systems involving chemical reactions.
1,2	Explain and use the terms: exothermic, endothermic, enthalpy of reaction ( ).
		Enthalpy change is heat transferred at constant pressure, a condition easily met in the laboratory. Only can be measured, not H for a system's initial or final state.
1,2	State the sign of for a reaction.
		Exothermic, heat released, is negative. Endothermic, heat absorbed, is positive.
1,2	Deduce from an enthalpy level diagram the relative stabilities of substances, and the sign of the enthalpy change for a conversion of one system to another.
		If the final state is more stable (lower on enthalpy level diagram), this implies that Hfinal < Hinitial and must be negative. Energy must be released in going to a more stable state.
1,2	Describe and explain the changes which take place at the molecular level in chemical reactions.
		Relate bond formation to the release of energy and bond breaking to the absorption of energy.
		Calculation of enthalpy changes (2h)
1,2	Calculate the energy transferred in changing the temperature of a pure substance.
1,2	Explain that enthalpy changes of reaction measure enthalpy changes for specific quantities of either reactants or products.
		Enthalpy changes are measured in kJ mol-1
1,2,3	Calculate the enthalpy change for a reaction in aqueous solution from experimental data on temperature changes and quantities of reactants.
		Enthalpy change of an acid-base reaction could be investigated. This could involve calculations of the enthalpy changes for a pure substance and for the quantity of reacting substance to obtain a molar enthalpy change of reaction.
1,2,3	Suggest suitable experimental procedures for measuring enthalpy changes of reactions in aqueous solution.
		Different reactions and variations in method could be explored, operating at constant pressure (open containers). Use of the bomb calorimeter is not required.
1,2,3	Describe, evaluate and explain the results of experiments on enthalpy changes.
		Use results from the experiments described in 6.2.3 and 6.2.4.
		Hess' law (1.5h)
1,2,3	Calculate the enthalpy change of a reaction which is the sum of two or more reactions with known enthalpy changes of reaction.
		Examples of simple two- and three-step processes should be used. Students should be able to construct simple enthalpy cycles. Cross reference to the combustion of hydrocarbons (topic 11). Students should be able to use Hess' law, but will not be required to state it.
		Bond enthalpies (1.5h)
1	Define average bond enthalpy.
		Bond enthalpies are normally quoted for the gaseous state and should be recognised as average values. They are most useful when only a few bonds are made and broken. Cross reference to the combustion of hydrocarbons (topic 11).
1,2	Calculate the enthalpy change of a reaction using bond enthalpies.

		Standard enthalpy changes of reaction (1h)
1,2	Explain and use the terms 'standard state' and 'standard enthalpy change of formation'.
		Standard state defined as 101 kPa, 298 K.
2	Calculate the enthalpy change of a reaction using standard enthalpy changes of formation.
		Lattice enthalpy (2h)
1	Explain and use the term 'lattice enthalpy'.
		See glossary. The value of lattice enthalpy will be taken as positive, although some texts consider lattice enthalpy to be the formation of the ionic lattice, with a negative value of Hlattice. The sign of Hlattice will always reveal which process is assumed.
1,2,3	Calculate enthalpy changes using Born-Haber cycles.
1,2,3	Predict the effect of the relative sizes and charges of ions on the relative lattice enthalpies of different ionic compounds.
		Students should be aware that the relative value of the theoretical lattice enthalpy increases with higher ionic charge and smaller ionic radius, due to increased attractive forces.
		Entropy (2h)
1,2,3	State and recognise factors which increase disorder in a system.
		An increase in disorder can result from mixing of different types of particles, change of state (increased distance between particles), increased movement of particles, or increased number of particles. Where the number of particles in the gaseous state increases, this factor is usually greater than any other possible factor.
1,2,3	Predict whether the entropy change for a given reaction or process would be positive or negative.
		From a reaction equation, candidates should be able to recognise a single factor which affects the value of S, and predict its sign, e.g., formation of a gas.
1,2	Calculate the standard entropy change for a reaction using values of absolute entropies.
		Spontaneity of a reaction (2h)
1,2	Explain that the tendency of a reaction to occur depends upon the total entropy change of the universe.
		Neither nor alone can reliably be used to predict the feasibility of a reaction. The ultimate criterion for feasibility is universe = system + surroundings > 0. The more common expression is , which must be negative for a spontaneous process. Students do not need to know how to derive this equation, but must be able to apply the principle.
2	Predict whether a reaction or process will be spontaneous by using the sign of .
1	State that
1,2	Calculate for a reaction using the equation or by using values of the standard Gibbs free energy change of formation.
1,2,3	Predict the effect of a change in temperature on the spontaneity of a reaction, given standard entropy and enthalpy changes.
		Using the equation (a)     is always negative if is negative and is positive; (b)     is negative at higher temperatures if both and are positive (an endothermic reaction is spontaneous where is greater than ); (c)     is negative at lower temperatures if both and are negative (an exothermic reaction is not spontaneous if is greater than ); (d)     is always positive if is positive and .