Energetics

6.1Energy in chemical reactions

Every process in the universe, in living cells or test tubes, in the atmosphere or in water, involves a change of energy. Some processes release energy and others need it, but the total amount of energy in the universe stays the same. Energetics is the study of the energy changes that go with all types of physical and chemical changes. This unit studies the role of energy in chemical reactions and how the heat released or absorbed in a reaction is measured and calculated.

In a chemical reaction, energy in the form of heat is either evolved or absorbed. The reason is that old bonds are broken and new bonds are formed. Bond breaking always consumes energy and bond making always releases energy.

6.1.1Units of thermal energy

Thermal energy is the energy in an object or system due to the movement of its molecules and atoms. The unit of heat or thermal energy in the SI system is the joule (J).

Another common unit of heat is the calorie. It is defined as the heat or thermal energy required to raise the temperature of one gram of water from $14.5^{\circ }\mathrm{C}$ to $15.5^{\circ }\mathrm{C}$.

6.1.2Thermochemical reactionsOften examined

When a chemical change takes place, energy is exchanged between the system and its surroundings. Energy has many forms, such as heat, light and work.

Thermochemical reaction

A chemical reaction which proceeds with the evolution or absorption of heat.

Thermochemical equation

A balanced chemical equation which also shows the heat change of the reaction.

Thermochemistry

The branch of chemistry which deals with the heat or thermal energy changes in chemical reactions.

For example, the burning of carbon:

There are two types of thermochemical reactions: exothermic and endothermic reactions.

1. Exothermic reactions

A chemical reaction that proceeds with the evolution of heat is called an exothermic reaction. The chemical system transfers energy to the surroundings as the reactants are converted to products. The burning of fuels is a highly exothermic reaction; the energy released can heat a room, drive an engine or cook food. Examples:

2. Endothermic reactions

A chemical reaction that proceeds with the absorption of heat is called an endothermic reaction. Heat is transferred from the surroundings to the system. Examples:

6.1.3Heat of reaction

The amount of energy released or absorbed depends on the amounts of substances that react. For example, the more fuel burns, the more energy is released. So, when an energy change is reported, the amounts of the chemicals that produce it must be reported as well. For this purpose energy changes are expressed for molar quantities of reactants and products.

Heat of reaction

The amount of heat evolved or absorbed in a chemical reaction when the molar quantities of reactants and products are the same as shown in the chemical equation.

Standard enthalpy change, $\Delta H^{\circ }$

The heat of reaction measured at $25^{\circ }\mathrm{C}$ (298 K) and one atmosphere pressure.

Reading a thermochemical equation:

This equation shows that 1 mole of solid carbon (12 g) reacts with 1 mole of oxygen gas (32 g) to give 1 mole of $\mathrm{CO}{}_2$ gas (44 g) at $25^{\circ }\mathrm{C}$ and 1 atmosphere pressure, and $393.5\mathrm{kJ}$ of heat is evolved.

The book asks what this second equation shows. It shows that 1 mole of hydrogen gas reacts with 1 mole of iodine gas to give 2 moles of hydrogen iodide gas, and $53.8\mathrm{kJ}$ of heat is absorbed (the sign is positive).

6.1.4Relation between enthalpy change and heat of reaction

Most chemical reactions occur at constant pressure, so the heat change in these reactions can be equated to the change in enthalpy. The heat of a reaction is therefore the change in enthalpy, \Delta H: the difference between the enthalpies of the products (final state) and the enthalpies of the reactants (initial state).

6.1.5Standard states and standard enthalpy changes

The value of \Delta H changes with the conditions, so standardised values of \Delta H are used. These values are calculated when all the substances are in their standard states (standard conditions). The superscript ^\circ in \Delta H^\circ shows that the value is for standard conditions.

6.1.6Conditions for standard heat of reaction

The standard states are as follows.

The book now defines eight important enthalpies. Each one is a standard enthalpy change for a particular kind of change; they are treated one by one in the sections that follow.

6.1.6 (1)Standard enthalpy of reactionOften examined

Standard enthalpy of reaction, \Delta H^\circ_\mathrm{r}

The enthalpy change in a chemical reaction when the reactants and products are in their standard states and their molar quantities are the same as shown by the balanced chemical equation.

Examples:

6.1.6 (2)Standard enthalpy of formationOften examined

Standard enthalpy of formation, \Delta H^\circ_\mathrm{f}

The enthalpy change that accompanies the formation of one mole of a compound from its elements, with all substances in their standard states.

To write a thermochemical equation for the enthalpy of formation:

1. Write the elements as reactants and 1 mole of the compound as the product.
2. Show the standard states of all the substances.
3. Finally balance the atoms (fractions such as \tfrac12 are allowed so that the product stays at one mole).

Standard enthalpies of formation of some substances are given in Table 6.1.

6.1.6 (3)Standard enthalpy of combustionOften examined

Standard enthalpy of combustion, \Delta H^\circ_\mathrm{c}

The enthalpy change when one mole of a substance is completely burnt in excess of oxygen under standard conditions.

Standard enthalpies of combustion of some substances are given in Table 6.2.

To write a thermochemical equation for the enthalpy of combustion:

1. Write 1 mole of the element or compound and oxygen as reactants.
2. Write the oxides of the given element, or of the elements present in the compound, as products.
3. Show the standard states of all the substances.
4. Finally balance the atoms.

6.1.6 (4)Standard enthalpy of atomization

Standard enthalpy of atomization, \Delta H^\circ_\mathrm{at}

The enthalpy change when one mole of gaseous atoms is formed from the element under standard conditions.

Examples:

6.1.6 (5)Standard enthalpy of neutralizationOften examined

Standard enthalpy of neutralization, \Delta H^\circ_\mathrm{n}

The amount of heat evolved when one mole of \ce{H+} ions from an acid combine with one mole of \ce{OH-} ions from a base to form one mole of water under standard conditions.

Strong acids and strong bases ionise completely in their aqueous solutions. So when solutions of \ce{NaOH} and \ce{HCl} are mixed, the only change that occurs is the formation of water; the \ce{Na+} and \ce{Cl-} ions remain in solution. The heat of neutralization is therefore due to the formation of water from \ce{H+} and \ce{OH-} ions.

6.1.6 (6)Standard enthalpy of solution

Standard enthalpy of solution, \Delta H^\circ_\mathrm{sol}

The enthalpy change when one mole of a substance is dissolved in so much solvent that further dilution results in no detectable heat change, under standard conditions.

Examples:

The first dissolving is endothermic (the solution cools) and the second is exothermic (the solution warms).

6.1.6 (7)Standard enthalpy of first electron affinity

Standard enthalpy of first electron affinity, \Delta H^\circ_\mathrm{EA}

The energy released when one mole of gaseous atoms gains an electron each to form one mole of gaseous ions with a single negative charge, under standard conditions.

6.1.6 (8)Lattice energy

Lattice energy

The energy released when one mole of an ionic compound (solid) is formed from its constituent gaseous ions.

6.2Bond energyOften examined

When a chemical reaction occurs, old bonds are broken and new bonds are formed. Breaking bonds always requires energy, and forming a bond always releases energy.

Bond dissociation energy

The energy required when one mole of a particular bond breaks apart to form neutral atoms.

Bond energy

The amount of energy released when one mole of that bond forms from neutral atoms.

The difference between the energy needed to break the bonds and the energy released on forming the new bonds decides whether the reaction absorbs or releases energy. For any chemical reaction, the enthalpy change is the sum of the bond dissociation energies of the reactants minus the sum of the bond energies of the products.

Enthalpy change from bond energies

The same formula gives the enthalpy change of the reaction between hydrogen and oxygen to produce water.

Calculation of bond energy for the enthalpy of reaction

You can calculate the bond energy per mole of methane by breaking all the \ce{C-H} bonds:

6.2.1Variability in bond energies

Some bond energies are relatively well defined and precise. Others are approximate and can vary depending on the specific molecules involved and the environment in which the bond occurs.

The \ce{C-C} bond energy in ethane (\ce{C2H6}) can be considered approximate because it depends on the general structure of the molecule. Similarly, the energy of a \ce{C-H} bond can vary with the chemical environment and the molecule in question. Some bonds have different strengths in different compounds: the \ce{O-H} bond energy in methanol (\ce{CH3OH}) is different from that in water because of differences in molecular structure and chemical environment.

6.3Heat and temperature

Heat as a form of energy

When water boils in a cooking pot, the lid jumps in the air after some time. The mechanical energy that makes the lid jump comes from the heat energy of the burner flame. This shows that heat is an energy that can be converted into other forms.

The law of conservation of energy states that energy cannot be produced in any physical or chemical process; it can only be changed from one form to another. Different forms of energy include light, mechanical, chemical, thermal, nuclear and magnetic energy.

Heat is all the energy produced by the movement of the molecules of a substance. Heat is a form of energy. When a body absorbs heat it becomes hotter, and when it releases heat it becomes colder. The effect of heat on a body is indicated by its temperature, which refers to the intensity of the heat of a body. Heat is measured in units of energy: the SI unit is the joule (J); it is also measured in calories.

Temperature is a measure of the average kinetic energy of the molecules of a substance. It describes the hotness or coldness of a substance or object. The SI unit of temperature is the kelvin (K); temperature is also measured on the Fahrenheit (^\circ\mathrm{F}) and Celsius (^\circ\mathrm{C}) scales. The amount of heat absorbed by a body is directly proportional to the temperature change.

Heat capacity and specific heat capacity

The amount of heat absorbed (q) by a substance is proportional to its temperature change. The constant of proportionality is the heat capacity.

Heat capacity

The amount of heat required to raise the temperature of a given amount of a substance by 1 kelvin. Unit: joule per kelvin (\mathrm{J\,K^{-1}}).

Specific heat capacity, c

The amount of heat required to raise the temperature of one gram of a substance by 1 kelvin. Unit: joule per gram per kelvin (\mathrm{J\,g^{-1}\,K^{-1}}).

Knowing the specific heat capacity of a substance, its mass and the temperature change when it is heated or cooled, the heat absorbed or released can be found.

6.3.1Measurement of heat of reactionOften examined

Calorimeter

A device that measures heat flow.

Calorimeters measure the heat released from a system either at constant pressure or at constant volume, so there are two types of calorimeter.

Constant pressure calorimeter

In a constant pressure calorimeter the pressure of the system is fixed. We need a thermally insulated container with a thermometer and a stirrer. For most purposes a coffee cup calorimeter is used.

The heat gained by the solution is the heat released by the reaction, and the reaction happens at constant pressure, so the heat of neutralization (or other reaction) per mole is

where n is the number of moles of the reactant that is used up (or of the product formed). The minus sign is there because a temperature rise means the reaction released heat, so \Delta H is negative.

6.4Hess's lawOften examined

Enthalpy is a state function, so the enthalpy change of a chemical reaction in going from some initial state to some final state does not depend on the path followed by the reaction. In going from a particular set of reactants to a particular set of products, the change in enthalpy is the same whether the reaction takes place in one step or in many steps. This principle is known as Hess's law.

Hess's law

G. H. Hess stated this law in 1840. It states that the enthalpy change in a chemical reaction is the same whether the reaction takes place in a single step or in several steps.

Mathematically, the sum of the enthalpy changes round a complete cycle is zero:

Suppose a reactant A changes into the product B in one step, with enthalpy change \Delta H. Now suppose the change takes place in three steps: A to C, C to D, and finally D to B.

If \Delta H_1, \Delta H_2 and \Delta H_3 are the enthalpy changes of the steps, then according to Hess's law:

Why Hess's law is useful

Many compounds cannot be prepared directly from their elements, and some of them cannot be decomposed into their constituent elements, for example \ce{CCl4}. Some elements do not burn completely because a protective covering forms on their surface, such as \ce{Al} and \ce{B}. So the enthalpies of formation of \ce{CCl4}, \ce{Al2O3} and \ce{B2O3} cannot be determined directly in a calorimeter. Hess's law is particularly useful for finding the enthalpies of formation of such compounds.

6.4.1Enthalpies of reactions from enthalpies of formationOften examined

It is often convenient to calculate \Delta H^\circ for a reaction from the standard enthalpies of formation, \Delta H^\circ_\mathrm{f}, of the reactants and products. For a given reaction:

where coeff is the coefficient of the substance in the balanced equation.

6.4.2Calculation of enthalpy of combustion

The enthalpy of combustion, \Delta H^\circ_\mathrm{c}, is the heat evolved when 1 mole of a substance burns completely in oxygen at standard conditions. It can be calculated by the same general formula as in section 6.4.1, with the combustion equation:

6.5Born–Haber cycleOften examined

The Born–Haber cycle is a special application of Hess's law to binary ionic compounds (\ce{M+X-}). It helps us to calculate the lattice energies of binary ionic compounds.

Lattice energy

The change in energy that takes place when separated gaseous ions are packed together to form one mole of an ionic solid.

Lattice energy cannot be determined directly. However, it can be determined indirectly by means of the Born–Haber cycle. Consider sodium chloride, whose standard enthalpy of formation is \Delta H^\circ_\mathrm{f}=-411\,\mathrm{kJ\,mol^{-1}}. The formation reaction can be thought of as taking place in several steps, one of which is the formation of the lattice. This complete sequence of reactions is called a cycle (Figure 6.5).

The five steps for \ce{NaCl}

The first three steps are all endothermic (energy is taken in), so they are drawn upward in the figure.

6.5.1Factors affecting lattice energyOften examined

The size of the lattice energy is influenced by the charge density of an ion. Charge density is the ratio of the ionic charge to the ionic radius of the ion.

Two factors have a significant effect on lattice energy.

1. Ionic charge

As the magnitude of the ionic charge increases, the lattice energy also increases. Ions with higher charges have stronger electrostatic attractions, and they release more energy when they come together to form an ionic lattice. For example, \ce{MgO} has a higher lattice energy than \ce{NaCl}, because \ce{Mg^2+} and \ce{O^2-} ions have a larger charge than \ce{Na+} and \ce{Cl-} ions.

2. Ionic radius

As the ionic radius increases, the lattice energy decreases. Larger ions keep the opposite ions at a greater distance, giving a weaker electrostatic attraction between the ions. Larger ions therefore need less energy to separate, which means a lower lattice energy. For example, \ce{LiF} and \ce{KF} have the same ionic charge but different ionic radii. The \ce{Li+} ion is smaller than the \ce{K+} ion, so \ce{LiF} has the stronger electrostatic attraction between its ions. So \ce{LiF} has a higher lattice energy than \ce{KF}.

6.5.2Role of hydration in the dissolving process

The energy change when a solution forms depends on three types of interaction: between solute and solute particles, between solvent and solvent particles, and between solute and solvent particles.

The strength of these two types of effect decides whether the dissolving is endothermic or exothermic. When a solute dissolves in water, the attractive forces between the solute particles must be overcome to break them out of the solid lattice. A strong force must bind these separate solute particles to the water molecules in solution. So the heat of solution measures the net flow of energy when a substance dissolves.

The energy needed to break the solute particles apart is the lattice energy (\Delta H_\text{lattice}) of the solute. The energy released when solute particles bind to water molecules is the enthalpy of hydration (\Delta H_\text{hyd}); for other solvents this energy is called the heat of dissolution. So the enthalpy of solution is the difference between the lattice energy and the enthalpy of hydration:

Endothermic and exothermic solution

If \Delta H_\text{solution} is very large and positive, the solute is unlikely to be soluble. If \Delta H_\text{solution} is moderately positive, the solute dissolves.

When ammonium nitrate dissolves in water, 25.7\,\mathrm{kJ\,mol^{-1}} of energy is absorbed and the flask becomes cold. This is because the crystal binds the \ce{NH4+} and \ce{NO3-} ions more tightly than the solution does.

On the other hand, if the solvent binds the solute particles more tightly than the crystal does, energy is released as the solute dissolves and the solution process is exothermic. When \ce{NaOH} dissolves in water, 44.5\,\mathrm{kJ\,mol^{-1}} of energy is released and the flask becomes hot.

So the solution process can be endothermic or exothermic depending on the difference between the lattice energy of the solute and the heat of hydration of its particles. Other examples are:

6.5.3Energy cycle involving enthalpy of solution, lattice energy and enthalpy of hydrationOften examined

A solid dissolving can be drawn as a cycle. The solid \ce{MX(s)} can go to the aqueous ions \ce{M+(aq) + X-(aq)} directly (the enthalpy of solution), or in two steps: first the lattice is broken into gaseous ions (the reverse of the lattice formation), and then the gaseous ions are hydrated.

According to Hess's law:

6.5.4Factors affecting enthalpy of hydration

The size of the enthalpy of hydration is influenced by the charge density of an ion, which is the ratio of the ionic charge to the ionic radius. So two factors affect the enthalpy of hydration.

1. Ionic charge

The charge of an ion affects the electrostatic interaction between the ion and the water molecules. Ions with a higher charge attract water molecules more strongly (ion–dipole interactions), so the enthalpy of hydration is larger: as the ionic charge increases, the hydration process becomes more exothermic. For example, the hydration enthalpy of \ce{MgO} is more exothermic than that of \ce{NaCl}, because \ce{Mg^2+} and \ce{O^2-} have a greater charge than \ce{Na+} and \ce{Cl-}.

2. Ionic radius

Smaller ions have a higher charge density and attract water molecules more strongly, so more energy is released when they hydrate and the enthalpy of hydration is more exothermic. Larger ions spread their charge out and attract water molecules less strongly, so less energy is released. For example, the hydration enthalpy of \ce{MgSO4} is greater than that of \ce{CaSO4}. The sulphate ion is common to both, so the difference must be due to \ce{Mg^2+} and \ce{Ca^2+}. Both are in group 2 and magnesium is above calcium, so \ce{Mg^2+} is smaller than \ce{Ca^2+}. Water molecules are attracted more strongly to \ce{Mg^2+}, so the hydration of \ce{MgSO4} is more exothermic than that of \ce{CaSO4}.

6.5.5Factors affecting electron affinity

Three factors affect electron affinity: nuclear charge, atomic size and shielding effect. The book gives the three names here and refers to section 10.5.1 for the details.

6.6EntropyOften examined

Which state of matter has more disorder, random order and minimum particle energy, solid or gas? The degree of disorder and randomness of a system is called its entropy. It is marked with the symbol S. Entropy is also described as the number of possible arrangements of the particles and their energy in a given system.

As a system moves from a more ordered state to a less ordered state, the entropy of the system tends to increase. Melting ice is a good example. In ice the molecules have a fixed and ordered arrangement and minimum energy. As the ice melts, the energy of its molecules increases, they gain more freedom of movement and become more disordered, so the entropy of the system increases. On the other hand, when water vapour condenses into a liquid, the molecules become more ordered and less energetic, so the entropy of the system decreases.

The entropy change of the system is denoted by \Delta S.

• If the entropy of the system increases, this is a positive change: \Delta S>0.
• If the entropy of the system decreases, this is a negative change: \Delta S<0.
• In general, when the temperature of a system increases, its entropy also increases, \Delta S>0.
• As the temperature of the system decreases, \Delta S<0.

6.6.1Calculation of the entropy change

The entropy change for a chemical reaction is calculated from the standard entropies, S^\circ, of the products and reactants:

6.7Gibbs free energyOften examined

Gibbs free energy is a measure of the amount of usable energy of a system at constant temperature and pressure. It is defined as the enthalpy of the system minus the product of its temperature and entropy. It is also known as the 'available energy'. It is denoted by the symbol G, after its discoverer, the American scientist Josiah Willard Gibbs, who introduced it in 1876.

Gibbs free energy is a state function: it does not depend on the path followed by a reaction. So the change in Gibbs free energy is given by

When the reaction is carried out at constant temperature, \Delta T=0, so \Delta(TS)=T\,\Delta S and

This equation is called the Gibbs–Helmholtz equation. \Delta G^\circ is particularly useful in deciding whether a chemical reaction is spontaneous or at equilibrium.

The sign and size of \Delta G^\circ show whether the process is spontaneous or at equilibrium. There are three possibilities.

Gibbs free energy is also useful for understanding phase transitions, such as from liquid to solid or gas and back. Consider the melting of ice:

At the melting point, \Delta G^\circ=0, so the solid and liquid phases are in equilibrium.

6.7Enthalpy and calorie contents of food

The amount of energy in a food or drink is measured in calories. The calories in food refer to the amount of energy stored in the food, which is released through the metabolic processes in our body. The enthalpy change linked with these processes corresponds to the amount of energy released or absorbed during digestion and metabolism.

The energy content of food is measured in kilocalories, and is often just called calories. One kilocalorie (kcal) is approximately equal to 4.184\,\mathrm{kJ}. This conversion factor relates the energy released in a chemical reaction to the calorie content of food.

Formula sheet

Key terms

Energetics

The study of the energy changes associated with all types of physical and chemical changes.

Thermochemistry

The branch of chemistry which deals with the heat or thermal energy changes in chemical reactions.

Thermal energy

The energy in an object or system due to the movement of its molecules and atoms.

Calorie

The heat required to raise the temperature of one gram of water from 14.5^\circ\mathrm{C} to 15.5^\circ\mathrm{C}; 1\,\text{cal}=4.18\,\mathrm{J}.

Thermochemical equation

A balanced chemical equation which also shows the heat change of the reaction.

Exothermic reaction

A reaction that proceeds with the evolution of heat; \Delta H is negative.

Endothermic reaction

A reaction that proceeds with the absorption of heat; \Delta H is positive.

Heat of reaction

The heat evolved or absorbed in a reaction when the molar quantities of reactants and products are the same as in the chemical equation.

Enthalpy change, \Delta H

The heat of reaction at constant pressure: H_\text{products}-H_\text{reactants}.

Thermodynamics

The study of all types of energy changes associated with chemical and physical changes.

State function

A property determined by the state of the system regardless of how that state was reached.

Internal energy

The sum of all kinds of energies of the particles of a system.

Enthalpy, H

The internal energy of a system plus the product of its pressure and volume: H=E+PV.

Standard states

1 atm for a gas; the most stable physical state at 1\,\mathrm{atm} and 25^\circ\mathrm{C} for an element or compound; 1\,\mathrm{M} for a substance in aqueous solution.

Standard enthalpy of reaction

The enthalpy change when reactants and products are in their standard states and in the molar quantities of the balanced equation.

Standard enthalpy of formation

The enthalpy change when one mole of a compound is formed from its elements, all in their standard states.

Standard enthalpy of combustion

The enthalpy change when one mole of a substance is completely burnt in excess oxygen under standard conditions.

Standard enthalpy of atomization

The enthalpy change when one mole of gaseous atoms is formed from the element under standard conditions.

Standard enthalpy of neutralization

The heat evolved when one mole of \ce{H+} from an acid combines with one mole of \ce{OH-} from a base to form one mole of water.

Standard enthalpy of solution

The enthalpy change when one mole of a substance is dissolved in so much solvent that further dilution gives no detectable heat change.

First electron affinity

The energy released when one mole of gaseous atoms gains one electron each to form gaseous ions with a single negative charge.

Lattice energy

The energy released when one mole of an ionic solid is formed from its gaseous ions.

Bond dissociation energy

The energy required to break one mole of a particular bond to form neutral atoms.

Bond energy

The energy released when one mole of a particular bond forms from neutral atoms.

Heat capacity

The heat required to raise the temperature of a given amount of a substance by 1 kelvin (\mathrm{J\,K^{-1}}).

Specific heat capacity

The heat required to raise the temperature of one gram of a substance by 1 kelvin (\mathrm{J\,g^{-1}\,K^{-1}}).

Calorimeter

A device that measures heat flow, at constant pressure or at constant volume.

Hess's law

The enthalpy change in a chemical reaction is the same whether the reaction takes place in one step or in several steps.

Born–Haber cycle

An application of Hess's law to ionic compounds that is used to find lattice energies.

Charge density

The ratio of the ionic charge to the ionic radius of an ion.

Enthalpy of hydration

The energy released when gaseous ions bind to water molecules.

Entropy, S

The degree of disorder and randomness of a system; the number of possible arrangements of its particles and their energy.

Gibbs free energy, G

The usable energy of a system at constant temperature and pressure: G=H-TS.

Spontaneous process

A process with \Delta G^\circ<0, which goes ahead by itself in the forward direction.

FBISE 2025 and 2026 questions

Questions from the FBISE HSSC-I 2026 (1st Annual) and FBISE HSSC-I 2025 (1st Annual) papers that belong to this unit, answered the way an examiner expects.

Exam focus

1. Apply Hess's law to find an enthalpy of reaction or of formation (for example \ce{S + O2 -> SO2}, and \ce{CH3COOH} from combustion data). (numerical)
   A 7-mark Section C question in 2025 Set D Q.6 OR and 3 marks in 2026 Set B Q.2 (iv) and in Set D (same idea). They are book numerical questions 12 and 16.
   Go to answer
2. Draw and explain a Born–Haber cycle for an ionic solid, and use it to find the lattice energy. (long)
   Asked in 2025 Set B Q.2 (xiv) (draw the cycle for \ce{KCl}). The book's cycle for \ce{NaCl} and the exercise for \ce{SrF2} use the same five steps.
   Go to answer
3. Define enthalpy of atomization, lattice energy and first electron affinity with examples. (short)
   Asked in three of the four sets: 2026 Set B Q.2 (xi) OR, 2026 Set D Q.2 (xi) OR and 2025 Set D Q.2 (i) OR. Learn the definition and one equation with its sign for each.
   Go to answer
4. Explain the enthalpy of solution through lattice energy and enthalpy of hydration (why \ce{NH4NO3} is endothermic and \ce{CaCl2} exothermic). (short)
   Asked in 2025 Set B (MCQ 5, +786-783=+3\,\mathrm{kJ}, and Q.2 (vi)). It is also book long question 21 (\ce{SrF2}).
   Go to answer
5. Define entropy, predict the sign of \Delta S, and use \Delta G=\Delta H-T\Delta S. (short)
   Asked in 2025: Set D Q.2 (vii) (sign of \Delta S), Set D MCQ 5 (G=H-TS) and Set B Q.2 (xiii) OR (calculation of \Delta S^\circ). Book short question 20 and long question 19.
   Go to answer
6. Calculate a bond energy or the enthalpy change of a reaction from bond energies. (numerical)
   Asked in 2026 Set D Q.2 (iv) (the \ce{H-Cl} bond energy, book Example 6.3).
   Go to answer
7. Know that exothermic reactions and combustion have a negative \Delta H. (short)
   Two MCQs in 2026 (Set B MCQ 5 and Set D MCQ 5). Compare book exercise MCQ v on heat of combustion.
   Go to answer
8. Calorimetry numericals with q=mc\Delta T and \Delta H=-q/n (neutralization, bomb calorimeter, heat of combustion). (numerical)
   Not in the 2025 or 2026 papers, but book numerical questions 5, 7 and 10 use it, and the SLOs list it.
   Go to answer