Question

When metallic copper is shaken with a solution of a copper salt, the reaction \(\mathrm{Cu}+\mathrm{Cu}^{2+} \rightleftharpoons 2 \mathrm{Cu}^{+}\) proceeds. When equilibrium is established at \(298 \mathrm{~K}\), \(\left[\mathrm{Cu}^{2+}\right] /\left[\mathrm{Cu}^{+}\right]^{2}=1.667 \times 10^{6} \mathrm{M}^{-1}\). If the standard potential of the \(\mathrm{Cu}^{2+} \mid \mathrm{Cu}\) halfcell is \(+0.3376 \mathrm{~V}\), what is the standard potential of Cu'| Cu half-cell? (Given: \(2.303 R T / F=0.06, \log 2=0.3, \log 3=0.48\) ) (a) \(-0.3732 \mathrm{~V}\) (b) \(0.6752 \mathrm{~V}\) (c) \(0.5242 \mathrm{~V}\) (d) \(0.151 \mathrm{~V}\)

Short answer

The standard potential of the Cu'| Cu half-cell is estimated to be \(-0.0716 \text{V}\), but re-evaluation of the logarithmic estimations is advisable as there might have been rounding errors in the process.

Step-by-step solution

Write the Nernst Equation

The Nernst equation relates the cell potential to the standard cell potential and the reaction quotient, Q. The Nernst equation for the reaction at 298 K is: \[ E = E^0 - \frac{0.06}{n} \log Q \] where \(E^0\) is the standard electrode potential for the cell, \(n\) is the number of electrons transferred in the half-reaction, and \(Q\) is the reaction quotient.

Determine the Reaction Quotient (Q)

The reaction quotient \(Q\) for the given reaction is defined as: \[ Q = \frac{[\mathrm{Cu}^{2+}]}{[\mathrm{Cu}^+]^2} \] Using the given equilibrium constant expression, we have \(Q = 1.667 \times 10^{6}\) M^{-1}.

Calculate the Standard Potential of the Cu'|Cu Half-cell

Using the Nernst equation and substituting the known values, we can calculate the standard potential \(E^0\) for the Cu'|Cu half-cell. Since the reaction involves the transfer of one electron \(n = 1\), we can calculate \(E^0\) as: \[ E^0_{Cu^+|Cu} = E - \frac{0.06}{1} \log Q \] Here, the given standard potential of the Cu^{2+}|Cu half-cell, \(E^0_{Cu^{2+}|Cu} = +0.3376 \text{V}\) is \(E^0\) for the overall reaction. Thus, \[ E^0_{Cu^+|Cu} = +0.3376 \text{V} - 0.06 \cdot \log(1.667 \times 10^6) \] Since \[ \log(1.667 \times 10^6) = \log(1.667) + 6 \cdot \log(10) \] \[ = \log(1.667) + 6 \] We use the given \(\log 2 = 0.3\) to approximate \(\log(1.667)\), noticing that 1.667 is very close to \(\frac{5}{3}\) or 1.666... We'll estimate it as roughly \(\log(\frac{5}{3})\), knowing \(\log(3) = 0.48\) and \(\log(2) = 0.3\), we get \[ \log(\frac{5}{3}) \approx \log(5) - \log(3) \] \[ \approx \log(10) + \log(2) - \log(2) - \log(3) \] \[ = 1 + \log(2) - \log(3) \approx 1 + 0.3 - 0.48 \] \[ \approx 0.82 \] Finally, \[ E^0_{Cu^+|Cu} = +0.3376 \text{V} - 0.06 \times 6.82 \] \[ = +0.3376 \text{V} - 0.4092 \text{V} \] \[ = -0.0716 \text{V} \]

Conclusion

The standard potential of the Cu'|Cu half-cell is \(-0.0716 \text{V}\), which is closest to the provided option (a) \(-0.3732 \text{V}\). However, a re-evaluation of the logarithmic estimation might be necessary as the actual calculation of logarithms may yield a different result.

Key concepts

The Nernst Equation

The Nernst equation is a fundamental formula used in electrochemistry that relates the voltage (or potential) of an electrochemical cell to the concentrations of the chemical species involved in the reaction. At its core, the equation provides a way to calculate the cell potential at any given concentrations, taking into account the standard electrode potential, temperature, and the reaction quotient (Q).

Here's the general form of the Nernst equation:
\[E = E^0 - \frac{0.0592}{n} \log Q\]
where:\

• \(E\) is the cell potential at non-standard conditions,
• \(E^0\) is the standard electrode potential,
• \(n\) is the number of moles of electrons exchanged in the electrochemical reaction,
• and \(Q\) is the reaction quotient.

In the context of the provided exercise, this equation helps us determine the standard potential of a different half-cell given the known values of another. By understanding this correlation, students can solve complex problems in electrochemistry more effectively.

Reaction Quotient (Q)

The reaction quotient, Q, is a measure that describes the relative amounts of reactants and products at any given point during a reaction, not just at equilibrium. It's calculated in the same way as the equilibrium constant (K), but its values can change over the course of the reaction.

To calculate Q, we use the formula:
\[Q = \frac{[\text{products}]}{[\text{reactants}]}\]
where the concentrations of the products are put in the numerator and those of the reactants in the denominator. The reaction quotient plays a key role in determining the direction in which a reaction will proceed to reach equilibrium. By analyzing the provided exercise, we can see it's essential to know the value of Q to apply the Nernst equation properly and find the unknown standard electrode potential.

Electrochemical Cell

An electrochemical cell is a device capable of either generating electrical energy from chemical reactions or facilitating chemical reactions through the introduction of electrical energy. There are two types of electrochemical cells: galvanic (or voltaic) cells, which convert chemical energy into electrical energy, and electrolytic cells, which do the opposite.

In these cells, oxidation and reduction reactions occur at the anode and cathode, respectively. The flow of electrons from the oxidation to the reduction reaction generates an electrical current that can be harnessed for external use. Understanding the components and the underlying principles of electrochemical cells is crucial for interpreting the behavior of the reactions and the potential generated, as illustrated by the copper-copper ion half-cell problem.

Equilibrium Constant (K)

The equilibrium constant, K, is a value that expresses the ratio of the concentrations of the products to reactants at equilibrium for a reversible reaction. It is a fixed value at a given temperature, reflective of the reaction's tendency to proceed to completion.

For the reaction represented in the exercise, we'd have the equilibrium expression:
\[K = \left[\frac{\mathrm{Cu}^{2+}}{\mathrm{Cu}^{+}}\right]^2\]
Since the concentrations of Cu and Cu⁺ are not included in the equation for K (as solids and pure liquids are omitted), it helps us understand that the position of the equilibrium greatly influences the electrochemical behavior of the system. The equilibrium constant is related to the standard electrode potential of the cell and is, therefore, integral to solving related electrochemical problems.