Depict the galvanic cell in which the reaction $Zn(s) + 2Ag^+(aq) \rightarrow Zn^{2+}(aq) + 2Ag(s)$ takes place. Further show:
(i) Which of the electrode is negatively charged?
(ii) The carriers of the current in the cell.
(iii) Individual reaction at each electrode.

Question

Accepted Answer

> [!TLDR]
> The cell is represented as $Zn(s) | Zn^{2+}(aq) || Ag^+(aq) | Ag(s)$. The zinc electrode is negatively charged (anode). Electrons carry current in the external circuit, while ions carry current inside the cell. At the anode: $Zn(s) \rightarrow Zn^{2+}(aq) + 2e^-$; at the cathode: $Ag^+(aq) + e^- \rightarrow Ag(s)$.

This is a classic Daniell-type cell, but with silver instead of copper. The key to understanding any galvanic cell is to see it as a device that separates the oxidation and reduction half-reactions, forcing electrons to travel through an external wire. That flow of electrons is what we harness as electrical energy.

The reaction given is spontaneous — zinc metal will naturally reduce silver ions because zinc is higher up in the electrochemical series (more reactive). Let's break down exactly how this works.

1. **Identify the half-reactions.**  
   The overall reaction is:  
   $$Zn(s) + 2Ag^+(aq) \rightarrow Zn^{2+}(aq) + 2Ag(s)$$  
   Zinc goes from oxidation state 0 to +2 — it loses electrons. Silver goes from +1 to 0 — it gains electrons. So:  
   - Oxidation (loss of electrons): $Zn(s) \rightarrow Zn^{2+}(aq) + 2e^-$  
   - Reduction (gain of electrons): $Ag^+(aq) + e^- \rightarrow Ag(s)$  
   Notice the reduction half-reaction needs only one electron, but the oxidation produces two. So we multiply the reduction half-reaction by 2 to balance electrons: $2Ag^+(aq) + 2e^- \rightarrow 2Ag(s)$.

2. **Which electrode is which?**  
   In a galvanic cell, the electrode where oxidation occurs is called the **anode**. The electrode where reduction occurs is the **cathode**.  
   - Anode: zinc metal (Zn) — it oxidises to $Zn^{2+}$ and releases electrons.  
   - Cathode: silver metal (Ag) — it is the surface where $Ag^+$ ions from solution gain electrons and deposit as solid silver.

3. **Which electrode is negatively charged?**  
   At the anode, zinc atoms lose electrons. These electrons build up on the zinc electrode, giving it a **negative char

Concept	Physical Meaning
$E^\circ_{\text{cell}}$	Voltage when all species are at 1 M (standard state)
$-\frac{RT}{nF} \ln Q$	Correction factor — adjusts voltage for real concentrations
$Q$	Tells you how far the reaction is from equilibrium

Step	What to Do	Why
1	Write balanced half-reactions	Identify $n$ (electrons transferred)
2	Write overall reaction	Determine $Q$ form
3	Plug into Nernst	Corrects $E^\circ$ for real conditions
4	Use $\log_{10}$ at $25^\circ \text{C}$	$0.0592/n$ is exam standard

Component	Answer
Cell representation	$Zn(s) \
(i) Negatively charged electrode	Zinc (anode)
(ii) Current carriers	Inside: ions; Outside: electrons
(iii) Anode reaction	$Zn(s) \rightarrow Zn^{2+}(aq) + 2e^-$
(iii) Cathode reaction	$Ag^+(aq) + e^- \rightarrow Ag(s)$

Aspect	Correct Answer	Common Mistake
Cell representation	$Zn(s) \\| Zn^{2+}(aq) \\|\\| Ag^+(aq) \\| Ag(s)$	Reversed order
Negative electrode	Zn (anode)	Ag (cathode)
Current carriers	Electrons (external) + Ions (internal)	Only electrons
Anode reaction	$Zn \rightarrow Zn^{2+} + 2e^-$	$Zn^{2+} \rightarrow Zn$
Cathode reaction	$2Ag^+ + 2e^- \rightarrow 2Ag$	$Ag^+ + e^- \rightarrow Ag$ (unbalanced)

Q.Depict the galvanic cell in which the reaction $Zn(s) + 2Ag^+(aq) \rightarrow Zn^{2+}(aq) + 2Ag(s)$ takes place. Further show:

Concept understanding — Cell Representation Nernst Equation

Cell Representation and the Nernst Equation: From Intuition to Precision

The Intuition First

The Precise Statement

Cell Representation: How We Write It

Applying the Nernst Equation to a Cell Representation

A Worked Example

The Key Insight

Why this formula?

Cell Representation & the Nernst Equation: Why It Works

The Core Question

1. The Fundamental Link: Gibbs Free Energy & Cell Potential

2. The Chemical Side: $\Delta G$ Depends on Concentration

3. Combining Both Sides: The Derivation

4. The "Why" in Plain Terms

5. The Common Form (log base 10)

6. Cell Representation: How to Write $Q$

7. Exam-Ready Summary

Method: Standard Cell Representation (IUPAC Convention)

Steps

Final Summary

Common Mistakes & How to Avoid Them

Mistake 1: Writing the Cell Representation in the Wrong Order

Mistake 2: Confusing Which Electrode is Negatively Charged

Mistake 3: Forgetting to Mention Both Carriers of Current

Mistake 4: Writing Half-Reactions with Wrong Stoichiometry

Mistake 5: Mixing Up Oxidation and Reduction at Electrodes

Quick Summary Table

Q.Depict the galvanic cell in which the reaction Zn(s)+2Ag+(aq)→Zn2+(aq)+2Ag(s)Zn(s) + 2Ag^+(aq) \rightarrow Zn^{2+}(aq) + 2Ag(s)Zn(s)+2Ag+(aq)→Zn2+(aq)+2Ag(s) takes place. Further show:

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Q.Depict the galvanic cell in which the reaction $Zn(s) + 2Ag^+(aq) \rightarrow Zn^{2+}(aq) + 2Ag(s)$ takes place. Further show:

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