Free Thermochemistry and Energetics of Chemical Reactions MCQs with Answers

1. In an exothermic reaction the enthalpy change is

• A. negative, because the products have less enthalpy than the reactants
• B. positive, because heat is released
• C. always zero
• D. negative, because the products have more enthalpy than the reactants

Explanation: Energy leaves the system as heat, so the products end up at a lower enthalpy and the difference, products minus reactants, is negative. The sign convention is taken from the point of view of the system, which is why released heat carries a minus sign even though the surroundings get warmer. Combustion and neutralisation are the standard exothermic examples.

2. Which of the following is an endothermic process?

• A. Combustion of methane
• B. Photosynthesis
• C. Neutralisation of an acid by an alkali
• D. Condensation of steam

Explanation: Photosynthesis absorbs light energy to build glucose from carbon dioxide and water, so the products store more energy than the reactants and the enthalpy change is positive. Combustion and neutralisation both release heat, and condensation releases the latent heat that was absorbed during evaporation. A useful check is that an endothermic reaction feels cold and usually needs a continuous energy supply.

3. A reaction mixture in a beaker becomes noticeably cold. This shows that the reaction is

• A. exothermic, since heat leaves the system
• B. at equilibrium
• C. endothermic, since heat is absorbed from the surroundings
• D. catalysed

Explanation: The system draws energy from its surroundings, and since the beaker and the solution are those surroundings their temperature falls. Dissolving ammonium chloride or ammonium nitrate in water is the usual classroom demonstration, and it is the principle behind an instant cold pack. An exothermic reaction would make the beaker warm instead.

4. In thermodynamics, the system is defined as

• A. everything in the universe
• B. the container only
• C. the energy of the reaction
• D. the part of the universe under study, separated from the surroundings by a boundary

Explanation: The system is whatever the chemist chooses to study, usually the reacting substances, and everything else is the surroundings, with the boundary between them being real or imaginary. An isolated system exchanges neither matter nor energy, a closed system exchanges energy only and an open system exchanges both. A reaction in an open beaker is therefore an open system.

5. A state function is a property whose value depends on

• A. the path taken to reach the present state
• B. the present state of the system only, not on how it was reached
• C. the time taken for the change
• D. the catalyst used

Explanation: Enthalpy, internal energy, entropy, pressure, volume and temperature are all state functions, so the change in any of them depends only on the initial and final states. This path independence is precisely what makes Hess's law valid and useful. Work and heat, by contrast, are path functions, since their values do depend on how the change was carried out.

6. Internal energy of a system is the

• A. heat supplied to the system only
• B. work done by the system only
• C. total of all kinetic and potential energies of the particles in the system
• D. energy lost to the surroundings

Explanation: Internal energy sums the translational, rotational and vibrational energy of the particles together with the potential energy stored in the bonds and intermolecular forces. Its absolute value cannot be measured, but the change in it can, which is all thermodynamics requires. It is a state function, so the change depends only on the initial and final conditions.

7. The first law of thermodynamics is a statement of

• A. the conservation of energy
• B. the conservation of mass
• C. the increase of entropy
• D. the constancy of temperature

Explanation: Energy can be converted from one form to another but can neither be created nor destroyed, which is expressed as the change in internal energy equalling the heat added to the system plus the work done on it. It follows that a machine producing energy from nothing is impossible. The tendency of entropy to increase is the second law, an entirely separate statement.

8. For a reaction carried out at constant pressure, the heat exchanged is equal to the change in

• A. internal energy
• B. enthalpy
• C. entropy
• D. free energy

Explanation: Enthalpy was defined precisely so that its change measures the heat of a process at constant pressure, which is the condition of almost every reaction done in an open vessel. At constant volume, as in a bomb calorimeter, the heat measured is the change in internal energy instead, since no expansion work is done. This is why the two quantities differ slightly for reactions involving gases.

9. The standard enthalpy of formation of an element in its most stable form is

• A. always negative
• B. always positive
• C. zero by definition
• D. equal to its atomic mass

Explanation: Forming an element from itself involves no change, so the value is defined as zero and this provides the reference point from which all other enthalpies of formation are measured. The convention applies only to the most stable allotrope, so graphite is zero while diamond is not. Choosing an arbitrary zero is necessary because absolute enthalpies cannot be measured.

10. Hess's law states that the enthalpy change of a reaction

• A. depends on the number of steps taken
• B. is the same whether the reaction occurs in one step or several
• C. is always negative
• D. depends on the catalyst used

Explanation: Because enthalpy is a state function, the total change depends only on the initial and final states, so a reaction that cannot be measured directly can be calculated through a series of steps that can. This is how the enthalpy of formation of carbon monoxide is obtained, since burning carbon always gives some carbon dioxide as well. Hess's law is essentially the first law applied to chemical change.