---
title: "Electrode potential; Cu/Ag cell EMF"
subject: "Engineering Chemistry"
chapter: 1
key: "semester_2/engineering_chemistry:electrode-potential-cell-emf"
source: https://www.bctengineer.com/q/semester_2/engineering_chemistry/electrode-potential-cell-emf
site: BCT Engineering
---

# Electrode potential; Cu/Ag cell EMF

Engineering Chemistry · Chapter 1: Electro-chemistry and Buffer

- Asked in: 2073 Bhadra, 1+2+2 marks
- Video: [Cell Potential Problems - Electrochemistry](https://www.youtube.com/watch?v=UzkLP8segcs) — The Organic Chemistry Tutor
- Answer checked by Claude Opus 5.5 (high effort) on 2026-10-14
- Page: https://www.bctengineer.com/q/semester_2/engineering_chemistry/electrode-potential-cell-emf
- Whole chapter: https://www.bctengineer.com/md/semester_2/engineering_chemistry/chapter-1

## Question

Define electrode potential. How do you measure standard electrode potential of zinc electrode? Find the Emf of the following cell at 25°C.
Cu/Cu²⁺(0.2M)//Ag⁺(0.1M)/Ag
Given: E°Ag⁺/Ag = 0.80 V and E°Cu²⁺/Cu = −0.34 V

## Short answer

Electrode potential is the tendency at electrode–solution interface. Zn standard potential is measured by coupling with SHE; emf 0.76 V gives E° = -0.76 V. For the given cell, E° = 0.80 - 0.34 = 0.46 V; applying Nernst for 2e⁻ with Q = 0.2/(0.1)² = 20 gives E = 0.46 - 0.038 = 0.42 V.

## Answer

#### Electrode potential
The tendency of an electrode (metal) to lose or gain electrons when it is in contact with a solution of its own ions is called its **electrode potential**. It is the potential difference at the metal–solution interface. When ions are at 1 M, gases at 1 atm and temperature 25°C, it is called the standard electrode potential (E°).

#### Measuring standard electrode potential of zinc
A zinc rod dipped in 1 M ZnSO₄ is connected through a salt bridge to a standard hydrogen electrode (SHE, Pt, H₂ 1 atm | H⁺ 1 M, E° = 0).

```text
Zn | Zn²⁺ (1 M) || H⁺ (1 M) | H₂ (1 atm), Pt
```
Zn is oxidised (anode), H⁺ is reduced (cathode). The measured emf of the cell is 0.76 V, and since E°SHE = 0:
E°cell = 0 − E°(Zn²⁺/Zn) = 0.76 V, so **E°(Zn²⁺/Zn) = −0.76 V**.

#### EMF of Cu | Cu²⁺ (0.2 M) || Ag⁺ (0.1 M) | Ag
The given E°(Cu²⁺/Cu) = −0.34 V is taken as the oxidation potential of Cu (reduction potential of Cu²⁺/Cu is +0.34 V). Ag⁺/Ag = +0.80 V.

- Anode: Cu → Cu²⁺ + 2e⁻
- Cathode: 2Ag⁺ + 2e⁻ → 2Ag
- Cell reaction: Cu + 2Ag⁺ → Cu²⁺ + 2Ag   (n = 2)

```text
E°cell = 0.80 − 0.34 = 0.46 V
Q = [Cu²⁺]/[Ag⁺]² = 0.2/(0.1)² = 20,  log Q = 1.301
E = E° − (0.0591/2) log Q
  = 0.46 − 0.02955 × 1.301
  = 0.46 − 0.0384 = 0.4216 V
```
**Answer: EMF ≈ 0.42 V.**

## Check yourself

The answer's key points as multiple-choice questions, with the correct option and why.

1. How is the standard electrode potential of zinc measured?
   A. Zn weighed before and after dipping
   B. Zn in 1 M ZnSO4 coupled with SHE through a salt bridge and emf measured
   C. Zn coupled with Cu and the emf subtracted from 1.10
   D. Zn dipped in water and pH measured
   Answer: B. Zn in 1 M Zn2+ is coupled with SHE; since E°SHE = 0, the measured emf gives E° of Zn. (Key point: Measuring E° of zinc)

2. The given E°(Cu2+/Cu) = -0.34 V is treated as an oxidation potential. What is E°cell for Cu + 2Ag+ -> Cu2+ + 2Ag?
   A. 1.14 V
   B. 0.46 V
   C. -0.46 V
   D. 0.80 V
   Answer: B. Cu2+/Cu reduction potential is +0.34 V, so E°cell = 0.80 - 0.34 = 0.46 V. (Key point: Using the given data)

3. For [Cu2+] = 0.2 M and [Ag+] = 0.1 M, what is Q?
   A. 2
   B. 20
   C. 0.2
   D. 0.05
   Answer: B. Q = [Cu2+]/[Ag+]^2 = 0.2/0.01 = 20. (Key point: Reaction quotient)

4. E = 0.46 - (0.0591/2) log 20. What is E?
   A. 0.50 V
   B. 0.42 V
   C. 0.46 V
   D. 0.38 V
   Answer: B. 0.02955 x 1.301 = 0.0384, so E = 0.46 - 0.038 = 0.42 V. (Key point: Final EMF)

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