Consider the following reaction:
`Cu(s)+2Ag^(+)(aq) to 2Ag(s)+Cu^(2+)(aq)`
Depict the galvanic cell in which the given reaction takes place.

Options

A. `Cu^(2+)(aq)|Cu(s)||Ag^(+)(aq)|Ag(s)`
B. `Cu(s)|Cu^(2+)(aq)||Ag^(+)(aq)|Ag(s)`
C. `Ag^(+)(aq)|Ag(s)||Cu^(2+)(aq)|Cu(s)`
D. `Ag(s)|Ag^(+)(aq)||Cu^(2+)(aq)|Cu(s)`

Explanation

To depict the galvanic cell for the reaction \( \text{Cu}(s) + 2\text{Ag}^+(aq) \rightarrow 2\text{Ag}(s) + \text{Cu}^{2+}(aq) \), we need to follow these steps:

### Step 1: Identify Oxidation and Reduction
- **Oxidation** is the loss of electrons. In this reaction, copper (Cu) is oxidized from its elemental state (0) to \( \text{Cu}^{2+} \) (2+), losing 2 electrons.
- **Reduction** is the gain of electrons. Silver ions (\( \text{Ag}^+ \)) are reduced from a +1 oxidation state to solid silver (0), gaining 1 electron for each silver ion.

### Step 2: Write Half-Reactions
- **Oxidation half-reaction**:
\[ \text{Cu}(s) \rightarrow \text{Cu}^{2+}(aq) + 2e^- \]
- **Reduction half-reaction**:
\[ 2\text{Ag}^+(aq) + 2e^- \rightarrow 2\text{Ag}(s) \]

### Step 3: Construct the Galvanic Cell
1. **Anode (Oxidation Site)**: This is where oxidation occurs. Place copper (Cu) at the anode.
- **Anode Reaction**: \( \text{Cu}(s) \rightarrow \text{Cu}^{2+}(aq) + 2e^- \)

2. **Cathode (Reduction Site)**: This is where reduction occurs. Place silver ions (\( \text{Ag}^+ \)) at the cathode.
- **Cathode Reaction**: \( 2\text{Ag}^+(aq) + 2e^- \rightarrow 2\text{Ag}(s) \)

3. **Salt Bridge**: Connect the two half-cells with a salt bridge to maintain electrical neutrality.

### Step 4: Label the Components
- **Left Side (Anode)**:
- Copper electrode (Cu)
- Copper ions in solution (\( \text{Cu}^{2+} \))

- **Right Side (Cathode)**:
- Silver ions in solution (\( \text{Ag}^+ \))
- Silver electrode (Ag)

### Final Representation of the Galvanic Cell:
The galvanic cell can be represented as:
\[ \text{Cu}(s) | \text{Cu}^{2+}(aq) || \text{Ag}^+(aq) | \text{Ag}(s) \]