Question

For each of the following reactions, indicate which elements are oxidized and which are reduced: a. \(2 \mathrm{Al}+3 \mathrm{Cu}^{2+} \longrightarrow 2 \mathrm{Al}^{3+}+3 \mathrm{Cu}\) b. \(2 \mathrm{H}^{+}+\mathrm{Ni} \longrightarrow \mathrm{Ni}^{2+}+\mathrm{H}_{2}\) c. \(2 \mathrm{I}^{-}+\mathrm{Cl}_{2} \longrightarrow \mathrm{I}_{2}+2 \mathrm{Cl}^{-}\)

Short answer

a. Al is oxidized, Cu is reduced; b. Ni is oxidized, H is reduced; c. I is oxidized, Cl is reduced.

Step-by-step solution

Assign Oxidation States

First, we need to determine the oxidation states of the elements in the reactants and products. For the reaction a, Aluminum (Al) starts with an oxidation state of 0 and ends with +3, indicating it is oxidized. Copper (Cu) starts at +2 and ends at 0, indicating reduction. In reaction b, Hydrogen goes from +1 to 0, indicating reduction, and Nickel (Ni) goes from 0 to +2, indicating oxidation. In reaction c, Iodine (I) goes from -1 to 0 (oxidation), and Chlorine (Cl) from 0 to -1 (reduction).

Identify Oxidation Changes

Oxidation is the loss of electrons, or an increase in oxidation state. Reduction is the gain of electrons, or a decrease in oxidation state. For reaction a, Al goes from 0 to +3, loss of electrons, thus Al is oxidized. Cu goes from +2 to 0, gain of electrons, thus Cu is reduced. For reaction b, H goes from +1 to 0, gain of electrons, thus H is reduced. Ni goes from 0 to +2, loss of electrons, thus Ni is oxidized. For reaction c, I goes from -1 to 0, loss of electrons, thus I is oxidized. Cl goes from 0 to -1, gain of electrons, thus Cl is reduced.

Identify Oxidizing and Reducing Agents

The oxidizing agent is the species that gets reduced, while the reducing agent is the species that gets oxidized. In reaction a, Cu²⁺ is the oxidizing agent and Al is the reducing agent. In reaction b, H⁺ is the oxidizing agent and Ni is the reducing agent. In reaction c, Cl₂ is the oxidizing agent and I⁻ is the reducing agent.

Key concepts

Oxidation States

The concept of oxidation states helps chemists keep track of the electrons in compounds and reactions. In a chemical equation, the oxidation state signifies the charge of an atom if its bonds were completely ionic. Assigning oxidation states is governed by a set of rules, such as the fact that the oxidation state of a pure element is always zero and that the sum of oxidation states in a neutral compound must total zero.

Understanding the changes in oxidation states during reactions, as demonstrated in the textbook steps, reveals which elements are oxidized or reduced. This is pivotal in classifying and predicting the outcomes of redox reactions. For example, in the solution provided, we observed Aluminum (Al) increased from an oxidation state of 0 to +3, illustrating a loss of electrons, which means it underwent oxidation.

Oxidizing and Reducing Agents

Oxidizing and reducing agents are essential characters in the theater of redox reactions. The oxidizing agent is the substance that takes electrons from another substance; therefore, it gets reduced. Conversely, the reducing agent is the one that gives away electrons and thus, is oxidized. To identify these agents, look at the elements changing oxidation states.

For example, in the exercise provided, in reaction b, H⁺ ions act as the oxidizing agent because they gain electrons and are reduced to H₂. Nickel (Ni), on the other hand, loses electrons and is oxidized, making it the reducing agent. By clearly defining the role of each reactant, scientists can predict the direction of electron flow and the feasibility of a reaction.

Electron Transfer in Redox

Redox reactions are fundamentally about electron transfer. When one element gets oxidized, another must be reduced; this balance maintains the conservation of charge. Oxidation involves the loss of electrons, while reduction involves the gain of electrons.

Looking at the reactions from the exercise, in reaction c, electrons transfer from the iodine ions (I⁻) to the chlorine molecules (Cl₂). The iodine ions are oxidized, evidenced by their oxidation state shifting from -1 to 0, resulting in the formation of iodine (I₂). Meanwhile, chlorine molecules gain those electrons, reducing their oxidation state from 0 to -1, and forming chloride ions (Cl⁻). This electron transfer is the driving force behind redox reactions, playing a key role in biological processes, industrial applications, and beyond.