1. **State the problem:** We have a galvanic cell with zinc (Zn) as the anode and copper (Cu) as the cathode. The cell diagram is given as $$\text{Zn(s)} | \text{ZnSO}_4(\text{aq}) || \text{CuSO}_4(\text{aq}) | \text{Cu(s)}$$. 2. **Identify the half-reactions:** - At the anode (oxidation): $$\text{Zn(s)} \rightarrow \text{Zn}^{2+} + 2e^-$$ - At the cathode (reduction): $$\text{Cu}^{2+} + 2e^- \rightarrow \text{Cu(s)}$$ 3. **Explain electron flow:** Electrons flow from the zinc anode to the copper cathode through the external circuit (voltmeter). 4. **Role of salt bridge:** The salt bridge containing KCl maintains electrical neutrality by allowing ion flow between the two half-cells. 5. **Overall cell reaction:** Combining the half-reactions, the net reaction is: $$\text{Zn(s)} + \text{Cu}^{2+} \rightarrow \text{Zn}^{2+} + \text{Cu(s)}$$ 6. **Summary:** Zinc is oxidized, losing electrons, and copper ions are reduced, gaining electrons, generating an electric current measurable by the voltmeter. This explains the galvanic cell operation and electron flow clearly.