# Electrochemical cells

> CIE A-Level Chemistry · Unit 19: Further electrochemistry
> Source: https://www.owlsprep.com/study/cie-9701-u19-electrochemical-cells/

This sub-topic covers the structure and function of electrochemical (galvanic) cells, calculation of standard cell potential from standard electrode potentials, and predicting redox reaction feasibility from cell potential values for CIE A-Level.

**Prerequisites:** [Standard hydrogen electrode and standard electrode potentials](https://www.owlsprep.com/study/cie-9701-u19-standard-electrode-potentials/); Redox reactions and half-equations

## Learning objectives

- Draw and label components of electrochemical cells
- Write correct standard cell notation
- Calculate standard cell potential from standard electrode potentials
- Predict reaction feasibility under standard conditions

## Structure and Components of an Electrochemical Cell

An electrochemical (galvanic) cell converts chemical energy from spontaneous redox reactions into electrical energy. It consists of two separate half-cells connected by an external wire and a salt bridge.

**Electrochemical (Galvanic) Cell** — A device that generates electrical energy from spontaneous, exergonic redox reactions, with oxidation in one half-cell and reduction in the other.

*Example:* A standard zinc-copper cell used to demonstrate spontaneous redox

- Anode: Negative electrode where oxidation (loss of electrons) occurs
- Cathode: Positive electrode where reduction (gain of electrons) occurs
- Salt bridge: Soaked in inert electrolyte, connects half-cells to balance charge
- External circuit: Wires connect electrodes, allowing electron flow to a load

> **tip**
>
> Electrons always flow from anode (negative) to cathode (positive) through the external wire, never through the electrolyte or salt bridge. Ions flow through the salt bridge to balance charge.

**Worked example:** Draw and label all key components of a zinc-copper electrochemical cell with zinc as anode

1. Draw two separate beakers for the two half-cells
2. Add a solid zinc strip to the first beaker, label as negative anode, fill with 1 mol dm⁻³ Zn(NO₃)₂ electrolyte
3. Add a solid copper strip to the second beaker, label as positive cathode, fill with 1 mol dm⁻³ Cu(NO₃)₂ electrolyte
4. Connect the two electrodes with an external wire, insert a high-resistance voltmeter
5. Place a KNO₃-soaked salt bridge between the two beakers, each end immersed in electrolyte
6. Label flow: electrons from Zn anode to Cu cathode; anions to anode, cations to cathode

## Standard Cell Notation

CIE exams often require standard cell notation to represent an electrochemical cell instead of a full drawing. There are clear, fixed conventions for this notation that you must follow.

1. Single vertical line | = phase boundary between solid electrode and aqueous electrolyte
2. Double vertical line || = the salt bridge between two half-cells
3. Anode (oxidation) half-cell is always written on the left
4. Cathode (reduction) half-cell is always written on the right

**Worked example:** Write cell notation for the reaction $Zn(s) + Cu^{2+}(aq) \rightarrow Zn^{2+}(aq) + Cu(s)$

1. Identify anode half-reaction: Zn oxidises to Zn²⁺, so this goes on the left
2. Write anode side with phase boundary: $Zn(s) | Zn^{2+}(aq, 1\ mol\ dm^{-3})$
3. Add double line for salt bridge: $Zn(s) | Zn^{2+}(aq, 1\ mol\ dm^{-3}) ||$
4. Add cathode side (Cu²⁺ reduces to Cu) on the right
5. $$\text{Final notation: } Zn(s) | Zn^{2+}(aq, 1\ mol\ dm^{-3}) || Cu^{2+}(aq, 1\ mol\ dm^{-3}) | Cu(s)$$

> **tip**
>
> If you get a negative E°cell after calculation, you almost certainly swapped anode and cathode in your notation. Check the order first!

## Calculating Standard Cell Potential

Standard cell potential ($E^\circ_{cell}$) is the overall potential difference between the two half-cells, measured under standard conditions (1 mol dm⁻³, 298 K, 1 atm pressure).

**Standard Cell Potential** — The potential difference between two electrodes when measured under standard conditions with no current flowing

*Notation:* E^\circ_{cell}

*Example:* E°cell for a zinc-copper cell is +1.10 V

The most reliable formula for calculation (avoids sign errors) is:

$$E^\circ_{\text{cell}} = E^\circ_{\text{cathode (reduction)}} - E^\circ_{\text{anode (oxidation)}}$$

**Worked example:** Given $E^\circ(Zn^{2+}/Zn) = -0.76\ V$ and $E^\circ(Cu^{2+}/Cu) = +0.34\ V$, calculate E°cell for the zinc-copper cell

1. Confirm: Zn is anode (oxidation), Cu is cathode (reduction)
2. Substitute values into the formula:
3. $$E^\circ_{\text{cell}} = 0.34\ V - (-0.76\ V)$$
4. Simplify and calculate the final result:
5. $$E^\circ_{\text{cell}} = 0.34 + 0.76 = +1.10\ V$$

**Check your understanding**

Test your understanding below:

1. Calculate E°cell for a cell with Fe anode (E° = -0.44 V) and Ag cathode (E° = +0.80 V):

   - +1.24 V
   - +0.36 V
   - -0.36 V
   - -1.24 V

   *Why:* Correct: 0.80 - (-0.44) = +1.24 V

2. Where is the silver half-cell written in cell notation?

   - Left of the salt bridge
   - Right of the salt bridge
   - Either side
   - Depends on E°cell

   *Why:* Silver is the cathode (reduction), which is always written on the right of the salt bridge

## Predicting Reaction Feasibility

The sign of E°cell tells us whether a reaction is feasible (spontaneous) under standard conditions.

**Feasible Reaction** — A reaction that can occur spontaneously without external energy input, predicted from E°cell under standard conditions

*Example:* Reaction of zinc with copper(II) ions is feasible

The core rule: positive E°cell means the forward reaction is feasible under standard conditions. Negative E°cell means the forward reaction is not feasible, and the reverse reaction is feasible.

> **warning**
>
> This rule only applies to standard conditions. Non-standard conditions (different concentration/temperature) can change the cell potential and alter feasibility.

**Worked example:** Predict if $2Ag(s) + 2H^+(aq) \rightarrow 2Ag^+(aq) + H_2(g)$ is feasible under standard conditions, given $E^\circ(Ag^+/Ag) = +0.80\ V$, $E^\circ(H^+/H_2) = 0\ V$

1. Split into half-reactions:
2. $$\text{Oxidation (anode): } 2Ag(s) \rightarrow 2Ag^+(aq) + 2e^-$$
3. $$\text{Reduction (cathode): } 2H^+(aq) + 2e^- \rightarrow H_2(g)$$
4. Calculate E°cell:
5. $$E^\circ_{\text{cell}} = E^\circ_{\text{cathode}} - E^\circ_{\text{anode}} = 0 - 0.80 = -0.80\ V$$
6. Conclusion: E°cell is negative, so the forward reaction is not feasible under standard conditions.

## Common pitfalls

- **Wrong:** Writing the anode half-cell on the right and cathode on the left in cell notation
  - Why it fails: Swapping the order leads to a negative E°cell, which incorrectly predicts the reaction is not feasible
  - Correct: Always place the oxidation (anode) half-cell on the left, reduction (cathode) on the right of the salt bridge
- **Wrong:** Forgetting to flip signs when using the E reduction + E oxidation method
  - Why it fails: All E values in the data booklet are reduction potentials, so oxidation E is the negative of the given value
  - Correct: Use the $E^\circ_{cell} = E^\circ_{cathode} - E^\circ_{anode}$ formula to eliminate sign errors entirely
- **Wrong:** Claiming a reaction with negative E°cell can never occur under any conditions
  - Why it fails: Negative E°cell only means non-feasible under standard conditions, non-standard conditions can make it feasible
  - Correct: Always add the qualifier 'under standard conditions' when predicting feasibility from E°cell
- **Wrong:** Stating electrons flow through the salt bridge or ions flow through the external wire
  - Why it fails: Electrons can only travel through the metallic wire, ions balance charge through the salt bridge
  - Correct: Remember: electrons flow anode → cathode through external wire; ions flow through salt bridge

## Cheatsheet

| Concept | Rule/Formula | Key Convention |
| --- | --- | --- |
| Cell components | Anode = oxidation (negative); Cathode = reduction (positive) | Electrons flow anode → cathode (external wire) |
| Cell notation | Anode \| Anode electrolyte \|\| Cathode electrolyte \| Cathode | \| = phase boundary, \|\| = salt bridge, anode left |
| E°cell calculation | $E^\circ_{cell} = E^\circ_{cathode} - E^\circ_{anode}$ | All values are standard reduction potentials |
| Feasibility | Positive E°cell = forward reaction feasible | Only applies to standard conditions |

## What's next

Electrochemical cells are the foundation for all advanced electrochemistry topics in CIE A-Level Chemistry, including practical applications like batteries, fuel cells, and electrolysis, as well as non-standard potential calculations with the Nernst equation. Mastery of cell conventions and E°cell calculations is essential, as these concepts appear frequently in both multiple-choice and extended response questions. E°cell is also directly related to Gibbs free energy change, connecting electrochemistry to the energetics topic you learned earlier, and explaining why spontaneous reactions have positive E°cell. Next, you will build on this knowledge to explore how non-standard conditions change cell potential.

- [Applications of electrode potentials](https://www.owlsprep.com/study/cie-9701-u19-applications-of-electrode-potentials/)
- [Further reaction kinetics](https://www.owlsprep.com/study/cie-9701-u20-overview/)
- [Rate equations and order of reaction](https://www.owlsprep.com/study/cie-9701-u20-rate-equations-and-order-of/)

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