Electrochemistry & Galvanic Cells practice questions

Question 1

A student constructs a galvanic cell operating under standard conditions using the two half-reactions shown below.

$\text{Ag}^+\text{(aq)} + e^- \rightarrow \text{Ag(s)} \quad E^\circ = +0.80 \text{ V}$
$\text{Zn}^{2+}\text{(aq)} + 2 e^- \rightarrow \text{Zn(s)} \quad E^\circ = -0.76 \text{ V}$

Which of the following balanced net-ionic equations represents the overall spontaneous reaction occurring in the cell?

Accepted Answer

$\text{Zn(s)} + 2\text{Ag}^+\text{(aq)} \rightarrow \text{Zn}^{2+}\text{(aq)} + 2\text{Ag(s)}$

Answer

$\text{Zn(s)} + \text{Ag}^+\text{(aq)} \rightarrow \text{Zn}^{2+}\text{(aq)} + \text{Ag(s)}$

Answer

2$\text{Ag(s)} + \text{Zn}^{2+}\text{(aq)} \rightarrow 2\text{Ag}^+\text{(aq)} + \text{Zn(s)}$

Answer

$\text{Ag(s)} + \text{Zn}^{2+}\text{(aq)} \rightarrow \text{Ag}^+\text{(aq)} + \text{Zn(s)}$

Question 2

A student constructs a galvanic cell operating under standard conditions using silver and lead electrodes submerged in standard $1.0 \, \text{M}$ nitrate solutions. The relevant standard reduction potentials are given in the table below.

| Half-Reaction | $E^\circ \, \text{(V)}$ |
| :--- | :--- |
| $\text{Ag}^+(\text{aq}) + e^- \rightarrow \text{Ag}(\text{s})$ | $+0.80$ |
| $\text{Pb}^{2+}(\text{aq}) + 2e^- \rightarrow \text{Pb}(\text{s})$ | $-0.13$ |

What is the standard cell potential, $E^\circ_{\text{cell}}$, for the galvanic cell?

Accepted Answer

$+0.93\text{ V}$

Answer

$-0.93\text{ V}$

Answer

$+0.67\text{ V}$

Answer

$+1.73\text{ V}$

Question 3

An electrolytic cell containing a solution of $\text{Ni(NO}_3)_2\text{(aq)}$ is operated at a constant current of $9.65 \text{ A}$. What is the time required to deposit $5.87 \text{ g}$ of $\text{Ni(s)}$ at the cathode?

Accepted Answer

$2,000 \text{ s}$

Answer

$4,000 \text{ s}$

Answer

$9,650 \text{ s}$

Answer

$19,300 \text{ s}$

Question 4

A student constructs a galvanic cell using a zinc electrode immersed in a $1.0 \text{ M ZnSO}_4\text{(aq)}$ solution and a copper electrode immersed in a $1.0 \text{ M CuSO}_4\text{(aq)}$ solution. The two half-cells are separated by a porous glass frit that permits ion movement between the solutions, as shown in the diagram. As the reaction $\text{Zn(s)} + \text{Cu}^{2+}\text{(aq)} \rightarrow \text{Zn}^{2+}\text{(aq)} + \text{Cu(s)}$ proceeds, which of the following correctly describes the migration of ions through the porous glass frit and provides the correct justification?

Accepted Answer

$\text{SO}_4^{2-}\text{(aq)}$ ions migrate toward the anode to neutralize the buildup of positive charge from $\text{Zn}^{2+}\text{(aq)}$ ions generated during oxidation.

Answer

$\text{SO}_4^{2-}\text{(aq)}$ ions migrate toward the cathode to neutralize the buildup of positive charge at the cathode.

Answer

$\text{Zn}^{2+}\text{(aq)}$ ions migrate toward the anode to balance the loss of positive charge as zinc metal is oxidized.

Answer

$\text{Cu}^{2+}\text{(aq)}$ ions migrate toward the anode to prevent the deposition of copper metal on the surface of the anode.

Question 5

A student investigates whether copper metal can provide the electrons needed to recover nickel metal from a solution. The proposed reaction is

$\text{Ni}^{2+}\text{(aq)}+\text{Cu(s)}\rightarrow\text{Ni(s)}+\text{Cu}^{2+}\text{(aq)}$

The relevant standard reduction potentials are shown below.

| Reduction half-reaction | $E^\circ$ |
|---|---:|
| $\text{Ni}^{2+}\text{(aq)}+2e^-\rightarrow\text{Ni(s)}$ | $-0.25\text{ V}$ |
| $\text{Cu}^{2+}\text{(aq)}+2e^-\rightarrow\text{Cu(s)}$ | $+0.34\text{ V}$ |

The experiment is conducted at $25^\circ\text{C}$ with both aqueous ions initially at $1.0\text{ M}$. Which choice correctly indicates whether the proposed reaction is thermodynamically favorable under standard conditions and whether an external voltage source is required to sustain the reaction as written?

Accepted Answer

Thermodynamically favorable: No; external voltage required: Yes

Answer

Thermodynamically favorable: Yes; external voltage required: No

Answer

Thermodynamically favorable: Yes; external voltage required: Yes

Answer

Thermodynamically favorable: No; external voltage required: No

Question 6

A student evaluates zinc granules as a low-cost method for recovering copper from wastewater. The student places clean $\text{Zn(s)}$ in a solution containing $\text{Cu}^{2+}\text{(aq)}$. After several minutes, chemical analysis confirms that $\text{Cu(s)}$ has formed and $\text{Zn}^{2+}\text{(aq)}$ is present. The process is represented by the following equation.

$$
\text{Zn(s)}+\text{Cu}^{2+}\text{(aq)}\rightarrow\text{Zn}^{2+}\text{(aq)}+\text{Cu(s)}
$$

Which choice correctly states whether $\text{Zn(s)}$ displaces $\text{Cu}^{2+}\text{(aq)}$ and explains the electron-transfer reason?

Accepted Answer

It does. Zinc atoms are more readily oxidized than copper atoms, so electrons from $\text{Zn(s)}$ reduce $\text{Cu}^{2+}\text{(aq)}$ to $\text{Cu(s)}$.

Answer

It does not. The $\text{Cu}^{2+}\text{(aq)}$ ions have already lost electrons and therefore cannot act as an oxidizing agent.

Answer

It does not. Copper atoms are less readily oxidized than zinc atoms.

Answer

It does. The $\text{Zn(s)}$ atoms gain electrons to form $\text{Zn}^{2+}\text{(aq)}$, while $\text{Cu}^{2+}\text{(aq)}$ loses electrons to form $\text{Cu(s)}$.

Question 7

A student constructs a galvanic cell using a zinc electrode in a solution of $1.0\text{ M }\text{Zn(NO}_3)_2\text{(aq)}$ and a copper electrode in a solution of $1.0\text{ M }\text{Cu(NO}_3)_2\text{(aq)}$. The half-cells are connected by an external wire through an ammeter and a salt bridge containing $\text{KNO}_3\text{(aq)}$. The ammeter registers a steady positive current. If the student removes the salt bridge while keeping the wire connected, which of the following best describes the change in current and the explanation for that change?

Accepted Answer

The current drops to zero because ion migration between half-cells stops, leading to charge accumulation that prevents further electron flow through the wire.

Answer

The current increases because removing the salt bridge decreases the total internal electrical resistance of the cell.

Answer

The current drops to zero because electrons are no longer able to flow through the salt bridge to complete the electrical circuit.

Answer

The current remains unchanged because electron movement through the external wire is independent of ion movement in the aqueous solutions.

Question 8

A student constructs a galvanic cell at $298 \text{ K}$ based on the balanced reaction below:

$$\text{Ni}(s) + 2\text{Ag}^+(aq) \rightarrow \text{Ni}^{2+}(aq) + 2\text{Ag}(s)$$

Initially, the cell operates under standard conditions with $[\text{Ag}^+] = 1.0 \text{ M}$ and $[\text{Ni}^{2+}] = 1.0 \text{ M}$, giving a standard cell potential $E^\circ_{\text{cell}} = +1.06 \text{ V}$. Distilled water is added to the $\text{Ag}^+(aq)$ half-cell, lowering $[\text{Ag}^+]$ to $0.10 \text{ M}$ while $[\text{Ni}^{2+}]$ remains $1.0 \text{ M}$. Which of the following predicts the effect of this modification on the cell potential, $E_{\text{cell}}$, and provides the correct reasoning?

Accepted Answer

$E_{\text{cell}}$ will be less than $+1.06 \text{ V}$ because decreasing $[\text{Ag}^+]$ increases the reaction quotient $Q$ above $1$, which decreases the thermodynamic driving force for the forward reaction.

Answer

$E_{\text{cell}}$ will be greater than $+1.06 \text{ V}$ because decreasing $[\text{Ag}^+]$ causes the system to shift toward the products to replace the lost reactant ions.

Answer

$E_{\text{cell}}$ will be less than $+1.06 \text{ V}$ because decreasing $[\text{Ag}^+]$ lowers the standard cell potential $E^\circ_{\text{cell}}$.

Answer

$E_{\text{cell}}$ will remain equal to $+1.06 \text{ V}$ because cell potential depends only on the identity of the electrodes and temperature, not on ion concentrations.

Question 9

A student electrolyzes an aqueous solution of $\text{CuSO}_4\text{(aq)}$ to deposit copper metal onto a cathode. To deposit $0.635 \text{ g}$ of $\text{Cu(s)}$ in $965 \text{ s}$, what constant electrical current, in amperes, must be passed through the cell? (Faraday's constant = $96,500 \text{ C/mol } e^-$; molar mass of $\text{Cu} = 63.5 \text{ g/mol}$)

Accepted Answer

$2.00 \, \text{A}$

Answer

$0.500 \, \text{A}$

Answer

$1.00 \, \text{A}$

Answer

$4.00 \, \text{A}$

Question 10

A student models the recovery of copper from a battery-recycling solution by passing a constant current through an electrolytic cell. Copper is deposited at the cathode according to the following half-reaction.

$\text{Cu}^{2+}\text{(aq)}+2\text{e}^-\rightarrow\text{Cu(s)}$

The current is $1.93\text{ A}$ and is applied for $1.00\times10^3\text{ s}$. Assume that all the current is used to deposit copper. The Faraday constant is $9.65\times10^4\text{ C/mol }\text{e}^-$, and the molar mass of copper is $63.5\text{ g/mol}$. What mass of copper is deposited on the cathode?

Accepted Answer

$0.635\text{ g}$

Answer

$0.318\text{ g}$

Answer

$1.27\text{ g}$

Answer

$2.54\text{ g}$

Question 11

A student constructs a standard galvanic cell operating at $298\text{ K}$ based on the half-reactions represented below.

$\text{Zn}^{2+}\text{(aq)} + 2e^- \rightarrow \text{Zn(s)} \quad E^\circ = -0.76\text{ V}$
$\text{Cu}^{2+}\text{(aq)} + 2e^- \rightarrow \text{Cu(s)} \quad E^\circ = +0.34\text{ V}$

The overall reaction occurring in the cell is $\text{Zn(s)} + \text{Cu}^{2+}\text{(aq)} \rightarrow \text{Zn}^{2+}\text{(aq)} + \text{Cu(s)}$. What is the value of the standard Gibbs free energy change, $\Delta G^\circ$, for the reaction at $298\text{ K}$?

Accepted Answer

$-212\text{ kJ/mol}_{rxn}$

Answer

$-425\text{ kJ/mol}_{rxn}$

Answer

$-318\text{ kJ/mol}_{rxn}$

Answer

$+212\text{ kJ/mol}_{rxn}$

Question 12

At 25°C, a student evaluates an electrochemical process for recovering zinc from an aqueous recycling stream. The reaction for the process is

$\text{Zn}^{2+}\text{(aq)} + \text{Cu(s)} \rightarrow \text{Zn(s)} + \text{Cu}^{2+}\text{(aq)}$

The standard cell potential for the reaction as written is $-1.10\text{ V}$. Faraday's constant is $96{,}500\text{ C/mol e}^-$. What is the standard Gibbs free-energy change, $\Delta G^\circ$, for the reaction?

Accepted Answer

$+212\text{ kJ/mol}$

Answer

$-212\text{ kJ/mol}$

Answer

$+21.2\text{ kJ/mol}$

Answer

$+106\text{ kJ/mol}$

Question 13

A student assembles the following concentration cell at $25^\circ\text{C}$, using a $\text{KNO}_3\text{(aq)}$ salt bridge.

$\text{Ag(s)}\mid\text{Ag}^+\text{(aq)},\ 0.010\text{ M}\parallel\text{Ag}^+\text{(aq)},\ \text{variable concentration}\mid\text{Ag(s)}$

The positive terminal of a voltmeter is connected to the right electrode. The student varies the right-half-cell concentration and graphs the measured potential difference, $E_{\text{right}}-E_{\text{left}}$, as a function of $\log\!\left(\dfrac{[\text{Ag}^+]_{\text{right}}}{[\text{Ag}^+]_{\text{left}}}\right)$.

Which statement correctly interprets point P and predicts the spontaneous direction of electron flow through an external wire?

Accepted Answer

At P, $[\text{Ag}^+]_{\text{right}}=100[\text{Ag}^+]_{\text{left}}$, and electrons flow from left to right because reduction occurs in the more concentrated half-cell.

Answer

At P, $[\text{Ag}^+]_{\text{right}}=\dfrac{1}{100}[\text{Ag}^+]_{\text{left}}$, and electrons flow from right to left because oxidation occurs in the more dilute half-cell.

Answer

At P, $[\text{Ag}^+]_{\text{right}}=2[\text{Ag}^+]_{\text{left}}$, and electrons flow from left to right because reduction occurs in the more concentrated half-cell.

Answer

At P, $[\text{Ag}^+]_{\text{right}}=100[\text{Ag}^+]_{\text{left}}$, and electrons flow from right to left because reduction occurs in the more dilute half-cell.

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