Question

For the following reactions, identify the substance oxidized, the substance reduced, the oxidizing agent, and the reducing agent. (a) \(\mathrm{Cu}+2 \mathrm{H}_{2} \mathrm{SO}_{4} \longrightarrow \mathrm{CuSO}_{4}+\mathrm{SO}_{2}+2 \mathrm{H}_{2} \mathrm{O}\) (b) \(3 \mathrm{SO}_{2}+2 \mathrm{HNO}_{3}+2 \mathrm{H}_{2} \mathrm{O} \longrightarrow 3 \mathrm{H}_{2} \mathrm{SO}_{4}+2 \mathrm{NO}\) (c) \(5 \mathrm{H}_{2} \mathrm{SO}_{4}+4 \mathrm{Zn} \longrightarrow 4 \mathrm{ZnSO}_{4}+\mathrm{H}_{2} \mathrm{~S}+4 \mathrm{H}_{2} \mathrm{O}\) (d) \(\mathrm{I}_{2}+10 \mathrm{HNO}_{3} \longrightarrow 2 \mathrm{HIO}_{3}+10 \mathrm{NO}_{2}+4 \mathrm{H}_{2} \mathrm{O}\)

Short answer

(a) Oxidized: Cu, Reduced: S, Oxidizing agent: H2SO4, Reducing agent: Cu. (b) Oxidized: SO2, Reduced: N in HNO3, Oxidizing agent: HNO3, Reducing agent: SO2. (c) Oxidized: Zn, Reduced: H in H2SO4, Oxidizing agent: H2SO4, Reducing agent: Zn. (d) Oxidized: I2, Reduced: N in HNO3, Oxidizing agent: HNO3, Reducing agent: I2.

Step-by-step solution

- Analyzing Reaction (a)

Determine the changes in oxidation states for the reactants and products. Cu starts with a 0 oxidation state and ends with +2 in CuSO4. The S in H2SO4 goes from a +6 in H2SO4 to +4 in SO2, indicating sulfur has been reduced. Hence, Cu is oxidized, and S is reduced. The oxidizing agent (which is reduced) is H2SO4 and the reducing agent (which is oxidized) is Cu.

- Analyzing Reaction (b)

Determine the changes in oxidation states for the reactants and products. S in SO2 goes from an oxidation state of +4 to +6 in H2SO4, indicating oxidation. Nitrogen in HNO3 goes from a +5 oxidation state to +2 in NO, indicating reduction. Thus, SO2 is the substance oxidized, and HNO3 is the substance reduced. The oxidizing agent is HNO3 and the reducing agent is SO2.

- Analyzing Reaction (c)

Determine the changes in oxidation states for the reactants and products. Zn starts with a 0 oxidation state and ends with +2 in ZnSO4. The S in H2SO4 remains at +6 in ZnSO4. The H in H2SO4 goes from +1 in H2SO4 to -1 in H2S, indicating a gain of electrons (reduction). Hence, Zn is oxidized and H is reduced. The oxidizing agent is H2SO4 and the reducing agent is Zn.

- Analyzing Reaction (d)

Determine the changes in oxidation states for the reactants and products. I2 starts with a 0 oxidation state and ends with +5 in HIO3, which indicates oxidation. Nitrogen in HNO3 goes from a +5 oxidation state to +4 in NO2, indicating reduction. Thus, I2 is the substance oxidized, and HNO3 is the substance reduced. The oxidizing agent is HNO3, and the reducing agent is I2.

Key concepts

Oxidation and Reduction

Understanding oxidation and reduction is essential in the study of redox reactions. These two processes always occur simultaneously and are respectively referred to as redox reactions.

Oxidation is the process where an atom or molecule loses electrons, which increases its oxidation state. This loss of electrons can be thought of as an increase in the 'oxidizing' character of the substance, as it often creates a more positive charge. For example, in the reaction where pure copper reacts with sulfuric acid to form copper sulfate, copper (Cu) loses electrons and is oxidized.

On the other hand, reduction is the gain of electrons by an atom or molecule, which results in a decreased oxidation state. The substance that undergoes reduction will typically gain a negative charge or become less positive in the process. Using the same reaction, sulfur (S) within sulfuric acid gains electrons and is thus reduced.

It's important to remember that an easy way to recall these processes is by the mnemonic 'OIL RIG': Oxidation Is Loss, Reduction Is Gain of electrons.

Oxidizing and Reducing Agents

Closely linked to the concepts of oxidation and reduction are oxidizing and reducing agents. An oxidizing agent is a substance that promotes oxidation by taking electrons from another substance, and as a result, it is itself reduced. Hence, the oxidizing agent is the species that gets reduced in a chemical reaction.

In the reaction between copper and sulfuric acid, the sulfuric acid (H₂SO₄) acts as the oxidizing agent because it accepts electrons from copper, which means it gets reduced as it undergoes a change from sulfur having a +6 oxidation state to +4 in the product sulfur dioxide (SO₂).

Conversely, a reducing agent is a substance that promotes reduction by giving electrons to another substance, and thus, it is itself oxidized. The reducing agent loses electrons and increases in oxidation state. In our example, copper (Cu) serves as the reducing agent because it donates electrons to sulfuric acid, causing the copper to be oxidized from a 0 to a +2 oxidation state.

Oxidation States

The oxidation state, also known as oxidation number, is a concept that provides a way to keep track of electrons during a redox reaction. It is assigned to indicate the degree of oxidation or reduction of an atom in a chemical compound. The rules for assigning oxidation states include elements in their elemental form having an oxidation state of 0 and for compounds, the sum of the oxidation states of all atoms should equal the net charge of the compound.

Let's consider reaction (c) as an example. Zinc (Zn) has an initial oxidation state of 0 since it is in its elemental form, and upon reacting, its oxidation state becomes +2 in the zinc sulfate (ZnSO₄) product. Likewise, the sulfuric acid's hydrogen goes from an oxidation state of +1 to -1 when it forms hydrogen sulfide (H₂S), indicating a reduction. Oxidation states are invaluable in identifying how electrons are transferred in a redox reaction.

Balancing Redox Equations

Properly balancing redox equations is crucial in accurately representing the stoichiometry of a redox reaction. This involves ensuring both mass and charge are balanced on both sides of the equation. A common method for balancing redox equations is the half-reaction method, which separates the overall reaction into two half-reactions - one for oxidation and one for reduction.

Each half-reaction is balanced separately for mass and charge, often necessitating the addition of water molecules, hydrogen ions (in acidic solutions), or hydroxide ions (in basic solutions) to balance oxygen and hydrogen atoms, while electrons are added to balance charges. After balancing, the half-reactions are added back together to get the final balanced equation.

As an exercise improvement advice, practice identifying and balancing each half-reaction first. This sub-division simplifies the process and serves as a solid foundation for mastering the art of balancing complex redox equations.