Free Electrochemistry MCQs with Answers

1. Oxidation is best defined as

• A. loss of electrons and an increase in oxidation number
• B. gain of electrons and a decrease in oxidation number
• C. gain of hydrogen only
• D. loss of oxygen only

Explanation: The modern definition is in terms of electrons, so oxidation is loss and reduction is gain, remembered as OIL RIG. Definitions based on adding oxygen or removing hydrogen are older special cases that fail for reactions in which neither element is present. Every oxidation must be accompanied by a reduction, since the electrons lost by one species are gained by another.

2. The oxidation number of sulphur in sulphuric acid, H2SO4, is

• A. plus 2
• B. plus 6
• C. minus 2
• D. plus 4

Explanation: Hydrogen contributes plus 1 each and oxygen minus 2 each, so 2 plus x minus 8 equals zero and x is plus 6. The value plus 4 belongs to sulphur in sulphur dioxide and in sulphurous acid, which is why it is the closest distractor. Sulphur at plus 6 is already at its maximum oxidation state, so concentrated sulphuric acid can act as an oxidising agent but not as a reducing one.

3. In the reaction Zn + CuSO4 gives ZnSO4 + Cu, the substance oxidised is

• A. the sulphate ion
• B. copper ions
• C. zinc
• D. copper metal

Explanation: Zinc goes from an oxidation number of zero as the metal to plus 2 in the sulphate, so it has lost electrons and is oxidised, which also makes it the reducing agent. The copper ions gain those electrons, fall from plus 2 to zero and are reduced. The sulphate ion is a spectator, unchanged on both sides of the equation.

4. An oxidising agent is a substance that

• A. loses electrons and is itself oxidised
• B. gains electrons and is itself reduced
• C. neither gains nor loses electrons
• D. always contains oxygen

Explanation: An oxidising agent oxidises something else by taking its electrons, so it is itself reduced in the process, and the same reasoning in reverse defines a reducing agent. Potassium permanganate and acidified potassium dichromate are common oxidising agents that contain oxygen, but chlorine is a strong oxidising agent with none, which disposes of the last option.

5. The oxidation number of oxygen in hydrogen peroxide, H2O2, is

• A. minus 2
• B. zero
• C. minus 1
• D. plus 2

Explanation: Hydrogen is plus 1, so the two oxygens must total minus 2 between them, giving minus 1 each, which is the exception to the usual rule of minus 2 for oxygen. This intermediate value is why hydrogen peroxide can act both as an oxidising agent, going to minus 2, and as a reducing agent, going to zero. In compounds with fluorine, oxygen takes a positive oxidation number.

6. A reaction in which the same element is both oxidised and reduced is called

• A. a neutralisation reaction
• B. a displacement reaction
• C. a precipitation reaction
• D. a disproportionation reaction

Explanation: In disproportionation one species with an intermediate oxidation state changes into two products, one at a higher and one at a lower state, as when chlorine reacts with cold dilute alkali to give chloride and hypochlorite. The element must therefore have an accessible oxidation state on both sides of its starting value. This is why sodium, always plus 1 in compounds, cannot disproportionate.

7. When balancing a redox equation by the ion electron method, the two half equations must be multiplied so that

• A. the number of atoms of oxygen is the same in both
• B. the number of electrons lost equals the number gained
• C. the charges on both sides are zero
• D. the coefficients are all whole numbers greater than one

Explanation: Electrons cannot appear in the final equation, so the two half equations are scaled until the electrons cancel exactly on addition, which is simply a statement that charge is conserved. Oxygen and hydrogen are balanced separately by adding water and hydrogen ions in acidic solution. The total charge on each side need not be zero, only equal.

8. In balancing a redox equation in acidic medium, oxygen atoms are balanced by adding

• A. hydroxide ions
• B. oxygen gas
• C. water molecules to the side deficient in oxygen
• D. hydrogen gas

Explanation: Water supplies the missing oxygen, and hydrogen ions are then added to the opposite side to balance the hydrogen introduced, which is valid because both are abundant in an acidic solution. In alkaline medium hydroxide ions are used instead, since hydrogen ions cannot be present in quantity. Adding elemental oxygen or hydrogen would change the chemistry rather than balance the equation.

9. The electrode potential of the standard hydrogen electrode is assigned a value of

• A. 0.00 volts by convention
• B. 1.00 volt
• C. minus 1.00 volt
• D. 0.76 volts

Explanation: No single electrode potential can be measured in isolation, so the hydrogen electrode is defined as zero and every other potential is quoted relative to it. The value 0.76 volts is the standard potential of the zinc electrode measured against it, which is why it appears as a distractor. The conditions specified are 1 molar hydrogen ions, 1 atmosphere of hydrogen gas and 298 K.

10. The standard hydrogen electrode consists of hydrogen gas at 1 atmosphere bubbled over

• A. a zinc rod in zinc sulphate solution
• B. platinised platinum foil dipping in 1 molar hydrogen ion solution
• C. a copper rod in copper sulphate solution
• D. a carbon rod in pure water

Explanation: Platinum is used because it is chemically inert and does not itself take part in the reaction, while the finely divided platinum black coating provides a large surface that catalyses the equilibrium between hydrogen gas and hydrogen ions. The solution must be exactly 1 molar in hydrogen ions and the temperature 298 K for the potential to be standard. Any reactive metal would corrode and destroy the reference.