Predicting Reactions (AQA A-Level Chemistry): Revision Notes

5.4.3 Predicting Reactions

Using Standard Electrode Potentials ($E^\theta$) to Predict Redox Reactions

Standard electrode potentials ($E^\theta$) allow us to predict the direction of a redox reaction by comparing the tendency of substances to gain or lose electrons. The reaction proceeds in the direction that results in a positive overall cell potential ($E^\theta_{\text{cell}}$).

Steps to Predict the Feasible Direction of a Redox Reaction

Step 1: Identify the Half-Equations:

• Write the two half-equations for the species involved in the reaction, ensuring both are written as reduction reactions (gaining electrons). Step 2: Determine Electrode Potentials:

• Look up the $E^\theta$ values for each half-equation using an electrochemical series table. Step 3: Identify Oxidation and Reduction:

• The half-equation with the more negative $E^\theta$ value will undergo oxidation (written in reverse) and act as the anode.

• The half-equation with the less negative or more positive $E^\theta$ value will undergo reduction and act as the cathode. Step 4: Combine the Half-Equations:

• Write the full redox equation by combining the half-equations.

• Calculate the overall $E^\theta_{\text{cell}}$ by subtracting the $E^\theta$ value of the oxidation half-cell from the $E^\theta$ value of the reduction half-cell:

$$E^\theta_{\text{cell}} = E^\theta_{\text{reduction}} - E^\theta_{\text{oxidation}}$$

• If $E^\theta_{\text{cell}}$ is positive, the reaction is feasible in the predicted direction.

Investigating Conditions: Concentration and Temperature Effects

Adjusting the concentration of ions and temperature in a voltaic cell influences the cell potential (EMF) due to shifts in equilibrium and reaction rates. Here are the expected results for each of these mini-experiments, along with explanations for these observations.

Effect of Concentration on EMF

Experiment

• Set up a zinc-copper voltaic cell.
• Change the concentration of $\text{Zn}^{2+}$ or $\text{Cu}^{2+}$ ions in their respective solutions.
• For example, increase the concentration of $\text{Cu}^{2+}$ ions in the copper half-cell or decrease the concentration of $\text{Zn}^{2+}$ ions in the zinc half-cell.

Expected Results

• Increasing $\text{Cu}^{2+}$ concentration: This shifts the equilibrium in the copper half-cell towards reduction, increasing the cell's EMF.
• Decreasing $\text{Cu}^{2+}$ concentration: This reduces the likelihood of reduction at the copper electrode, lowering the EMF.
• Increasing $\text{Zn}^{2+}$ concentration: This promotes oxidation in the zinc half-cell, which will decrease the EMF.
• Decreasing $\text{Zn}^{2+}$ concentration: This lowers the tendency for zinc to oxidise, which will increase the EMF slightly.

Effect of Temperature on EMF

Experiment

• Use the zinc-copper voltaic cell setup and vary the temperature of the solutions.
• Measure the EMF at room temperature, then gradually heat the cell to higher temperatures and observe changes in the cell voltage.

Expected Results

• Increasing temperature: The EMF of the cell decreases if the overall cell reaction is exothermic (which is typical for voltaic cells).
• Decreasing temperature: The EMF of the cell increases if the reaction is exothermic.

Explanation

In an exothermic cell reaction, higher temperatures reduce EMF because, according to Le Chatelier's principle, the equilibrium position shifts to favour the endothermic reverse reaction to counteract the added heat. This shift reduces the cell potential. Lower temperatures favour the exothermic forwards reaction, which typically increases the EMF.

Overview of Observations

• Higher $\text{Cu}^{2+}$ concentration increases EMF; higher $\text{Zn}^{2+}$ concentration decreases it.
• Higher temperatures typically reduce EMF in exothermic reactions, while lower temperatures increase EMF.