All Free Chemistry MCQs with Answers

171. 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.

172. 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.

173. 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.

174. 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.

175. Given that C + O2 gives CO2 has an enthalpy change of minus 393 kJ per mole, and CO + half O2 gives CO2 has an enthalpy change of minus 283 kJ per mole, the enthalpy of formation of carbon monoxide is

• A. minus 676 kJ per mole
• B. plus 110 kJ per mole
• C. minus 110 kJ per mole
• D. minus 283 kJ per mole

Explanation: Forming carbon monoxide and then burning it must total the same as burning carbon directly, so the unknown equals minus 393 minus the value of minus 283, giving minus 110 kJ per mole. Adding the two values instead of subtracting gives minus 676, which is the trap. This calculation is the standard textbook illustration of Hess's law.

176. The enthalpy of neutralisation of a strong acid by a strong alkali is almost constant at about minus 57 kJ per mole because

• A. all acids contain the same number of hydrogen atoms
• B. the salts formed have similar masses
• C. the reaction is always carried out at the same concentration
• D. in every case the same reaction occurs, hydrogen ions combining with hydroxide ions to form water

Explanation: Strong acids and strong alkalis are fully ionised, so the spectator ions take no part and the only chemical change is the formation of water from its ions, which releases the same energy whatever salt is produced. With a weak acid the value is smaller, because some energy is absorbed in ionising the acid first. This constancy is strong evidence for the ionic theory of neutralisation.

177. The standard enthalpy of combustion is the enthalpy change when

• A. one mole of a substance is burnt completely in excess oxygen under standard conditions
• B. one gram of a substance is burnt
• C. one mole of a substance is formed from its elements
• D. one mole of a substance is dissolved in water

Explanation: The definition specifies one mole, complete combustion and excess oxygen, and combustion values are always negative because burning releases energy. Incomplete combustion would give a different and smaller value, which is why excess oxygen is stipulated. Formation from elements is a separate quantity that may be positive or negative.

178. Bond breaking and bond formation are respectively

• A. both exothermic
• B. both endothermic
• C. exothermic and endothermic
• D. endothermic and exothermic

Explanation: Energy must be supplied to pull bonded atoms apart, so breaking is endothermic, while forming a bond releases energy and is exothermic. Whether the overall reaction is exothermic depends on which total is larger, so a reaction releasing more energy in making new bonds than it spends breaking old ones is exothermic. This is the reasoning behind calculating enthalpy change from bond energies.

179. In a bomb calorimeter, the heat of reaction measured corresponds to the change in

• A. enthalpy, since the pressure is constant
• B. internal energy, since the volume is constant
• C. entropy
• D. free energy

Explanation: A bomb calorimeter is a sealed rigid vessel, so the volume cannot change, no expansion work is done and all the energy released appears as heat, which equals the change in internal energy. A reaction in an open vessel is at constant pressure and gives the enthalpy change instead. The two differ only by the work done pushing back the atmosphere.

180. When 1 mole of methane burns, the energy released is used to heat water. This illustrates that

• A. energy is created during combustion
• B. energy is destroyed during combustion
• C. energy is transferred from the system to the surroundings but the total remains constant
• D. the enthalpy of the system increases

Explanation: The chemical energy stored in the bonds of methane and oxygen is converted into heat, which passes to the water, so nothing is created or destroyed and the first law holds. The system loses enthalpy, which is why the value is negative, while the surroundings gain the same amount. Every combustion process is an example of this transfer.