Problem 83 - Chemistry and Chemical Reactivity

Question

A solution of KI is added dropwise to a pale blue solution of \(\mathrm{Cu}\left(\mathrm{NO}_{3}\right)_{2} .\) The solution changes to a brown color and a precipitate forms. In contrast, no change is observed if solutions of KCl and KBr are added to aqueous \(\mathrm{Cu}\left(\mathrm{NO}_{3}\right)_{2} .\) Consult the table of standard reduction potentials to explain the dissimilar results seen with the different halides. Write an equation for the reaction that occurs when solutions of KI and \(\mathrm{Cu}\left(\mathrm{NO}_{3}\right)_{2}\) are mixed.

Step-by-Step Solution

Answer

Cu(NO₃)₂ and KI react to form CuI precipitate and I₂, due to favorable redox potentials.

Step 1: Determine Possible Reactions

When a solution of KI is added to Cu(NO₃)₂, the KI can potentially react with any copper ions (Cu²⁺) present. However, in order to determine if and what reaction will occur, we need to consider the reduction and oxidation possibilities, specifically focusing on the reduction potential of Cu²⁺ being reduced to Cu⁺ or Cu, and I⁻ being oxidized to I₂.

Step 2: Check Standard Reduction Potentials

Reference the standard reduction potentials table. The reduction potential for Cu²⁺ to Cu⁺ is +0.15 V, and for Cu²⁺ to Cu is +0.34 V (though usually leads to a lower state like Cu⁺ here with iodine). The oxidation potential for I⁻ to I₂ is -0.54 V. Since I⁻ to I₂ implies oxidation, its potential is considered positive, +0.54 V, when written as a reduction half-reaction: I₂ + 2e⁻ → 2I⁻.

Step 3: Analyze Reaction Feasibility

The feasibility of this redox reaction follows from the difference in potentials. The reaction Cu²⁺ + 2I⁻ → CuI(s) + 0.5I₂ has a net potential obtained by adding the reduction potential of Cu²⁺ → Cu⁺ (+0.15 V) and the oxidation potential of I⁻ to I₂ (+0.54 V). The combined reaction is spontaneous with a positive net potential (0.15 + 0.54 = 0.69 V).

Step 4: Write the Reaction Equation

Given the feasible electron transfer: \[\text{Cu}^{2+}_{(aq)} + 2\text{I}^-_{(aq)} \rightarrow \text{CuI}_{(s)} + 0.5\text{I}_2 (aq)\]This results in the brown color and formation of copper(I) iodide precipitate. In contrast, KCl and KBr do not oxidize because Cl⁻ and Br⁻ have much lower oxidation potentials compared to I⁻ and thus do not lead to a spontaneous reaction with Cu²⁺.

Key Concepts

Understanding Standard Reduction Potentials

One of the most important concepts in predicting redox reactions is the standard reduction potential, which tells us how likely a substance is to gain electrons. Each half-reaction has a standard reduction potential value, typically given in volts (V).

When dealing with redox reactions, we consult these values to determine which substances can undergo oxidation and reduction. For example, the reduction of copper ions ( Cu^{2+} ) can proceed to either Cu^+ or Cu , each with different reduction potential values of +0.15 V and +0.34 V, respectively.

Meanwhile, the conversion of iodide ions ( I^- ) to iodine ( I_2 ) is another half-reaction, but it initially presents as an oxidation process, which means its potential is considered positive when evaluated as a reduction: +0.54 V.

• A positive reduction potential indicates a greater likelihood to be reduced.
• A more negative or positive oxidation potential, when inverted, suggests a substance tends to oxidize easier than others.

The Role of Copper Halides

Copper halides are compounds formed between copper and halogen atoms. CuI (copper iodide) is the result in our original exercise where potassium iodide (KI) is added to copper nitrate. It turns brown due to the precipitate formation of copper(I) iodide.

Each halide ion, like chloride ( Cl^- ), bromide ( Br^- ), or iodide ( I^- ), might behave differently when it meets copper ions. This influences the reaction's outcome because iodide has a higher tendency to be oxidized, compared to chloride or bromide.

This differential reactivity is attributed to the diverse standard oxidation potentials of these halides. Iodide's relatively high oxidation potential leads to the successful reduction of copper, thereby creating the visible precipitate and color change:

• Cl^- and Br^- have lower standard oxidation potentials than I^- .
• This explains why Cl^- and Br^- do not react with Cu^{2+} , keeping the solution unchanged.

Analyzing Reaction Feasibility

The feasibility of a chemical reaction, especially a redox process, is largely predicted based on overall cell potential. If a reaction's net potential is positive, it is likely to occur spontaneously.

In our example: Cu^{2+}_{(aq)} + 2I^-_{(aq)} → CuI_{(s)} + 0.5I_2(aq) , by calculating with the standard reduction potentials we get 0.15 ext{ V} + 0.54 ext{ V} = 0.69 ext{ V} , which is positive.

This positive value indicates that the reaction between copper ions and iodide is feasible and spontaneous under standard conditions. In contrast, because Cl^- and Br^- have lower oxidation potentials, their reactions with Cu^{2+} are not feasible, leading to no noticeable change.

• When standard potentials form a positive net value, expect the reaction to proceed.
• Negative or zero net potentials imply that the reaction is not likely to occur on its own under standard conditions.