Commercial Applications of Electrochemical Cells

Learning Objective 3.1.11.2 Understand how electrochemical cells are applied in practice, including non-rechargeable, rechargeable, and fuel cells, and evaluate their benefits and risks.

Quick Notes

Electrochemical cells are used as sources of electrical energy in batteries and fuel cells.
Three main types of cells:
- Non-rechargeable Cells – Single-use, irreversible reactions (e.g., alkaline batteries).
- Rechargeable Cells – Reversible reactions allow recharging (e.g., lithium-ion batteries).
- Fuel Cells – Continuous electricity generation using a fuel without the need for recharging (e.g., hydrogen-oxygen fuel cells).
Reactions you need to know:

Electrode reactions in a lithium-ion battery:
- Positive electrode: Li⁺ + CoO₂ + e⁻ → Li⁺[CoO₂]⁻
- Negative electrode: Li → Li⁺ + e⁻
Hydrogen-oxygen fuel cell (alkaline conditions):
- Anode: 2H₂ + 4OH⁻ → 4H₂O + 4e⁻
- Cathode: O₂ + 2H₂O + 4e⁻ → 4OH⁻
- Overall: 2H₂ + O₂ → 2H₂O

Full Notes

1. How Electrochemical Cells Generate Electricity

Electrochemical cells can be used to convert chemical energy into electrical energy through redox reactions. These cells are also referred to as Galvanic or voltaic cells.

When two different half-cells are connected together by a wire and salt bridge, electrons flow from the negative electrode (anode) to the positive electrode (cathode), generating an electric current.

Types of Electrochemical Cells

Non-Rechargeable Cells

These are single-use cells where the reactions generating the electrical current are not reversible. Once the reactants get used up, the potential difference in the cell decreases and it can no longer produce an electrical current. The cell (or battery) cannot be recharged.

For example:Zinc and carbon based cells are single use.

AQA A-Level Chemistry diagram of a non-rechargeable alkaline battery cell setup

Rechargeable Cells

These are cells in which the reactions generating the electrical current are reversible. As the reactants get used up and their concentrations decrease, the cells can be recharged (products turned back into reactants) by applying an external current.

A common example of a rechargeable cell is a lithium-ion cell, used to make lithium-ion batteries.

Reactions in a lithium-ion battery:

Positive electrode (cathode, where reduction happens):

AQA A-Level Chemistry lithium-ion battery positive electrode reduction Li+ + CoO2 forming Li+[CoO2]- reaction diagram

Negative electrode (anode, where oxidation happens):

AQA A-Level Chemistry lithium-ion battery negative electrode oxidation reaction Li to Li+ diagram

When the cell is in use, Li is oxidised, and CoO₂ is reduced. Recharging reverses the reaction.

Matt’s exam tip

If asked to give the overall cell reaction that occurs during recharging – it is just the opposite of the reaction that occurs spontaneously. The cell EMF (E°cell) of a cell being recharged will be a negative value (the cell EMF of a cell’s spontaneous reaction will always be positive).

Fuel Cells

Fuel cells generate electricity continuously as long as fuel is supplied meaning they don’t require recharging – instead, fuel is constantly replenished.

The most common example is the hydrogen-oxygen fuel cell. Hydrogen gas (H₂) is oxidised and oxygen gas (O₂) is reduced, with water (H₂O) being the overall product.

For Example:Hydrogen-Oxygen Fuel Cell (Alkaline Conditions)

AQA A-Level Chemistry hydrogen-oxygen alkaline fuel cell diagram showing electrode reactions and overall reaction

Anode (oxidation): 2H₂ + 4OH⁻ → 4H₂O + 4e⁻
Cathode (reduction): O₂ + 2H₂O + 4e⁻ → 4OH⁻
Overall: 2H₂ + O₂ → 2H₂O

Using Electrode Potentials to Deduce Cell Reactions

Given standard electrode potentials (E° values), we can determine:

Which species is oxidised and which is reduced.
The E_cell (EMF) of the cell using:

AQA A-Level Chemistry formula box showing E°cell = E°(cathode) − E°(anode)

Remember for a spontaneous cell reaction:

E°_cathode is the more positive electrode (reduction).
E°_anode is the more negative electrode (oxidation).

Worked Example

Lithium Cell

Positive Electrode (Reduction): Li⁺ + CoO₂ + e⁻ → Li⁺[CoO₂]⁻ (E° = +0.56 V)
Negative Electrode (Oxidation): Li → Li⁺ + e⁻ (E° = −3.04 V)

Cell EMF Calculation:
E°cell = (+0.56) − (−3.04) = +3.60 V

The positive EMF confirms the reaction is spontaneous.
The EMF when the cell is being recharged would be −3.60 V (negative because energy is being put in to the system to force the reactions to occur in the opposite direction).

Benefits and Risks of Using Electrochemical Cells

Advantages of Electrochemical Cells

Provide a portable source of electrical energy.
Rechargeable cells can be reused many times.
Fuel cells provide continuous supply as long as fuel is available.
Hydrogen fuel cells produce only water as waste product (no CO₂ emissions at point of use).

Disadvantages and Risks of Electrochemical Cells

Non-rechargeable cells create waste once spent.
Rechargeable cells degrade over time and lose efficiency.
Production of fuel cells and batteries requires toxic/rare materials (e.g., lithium, cobalt).
Hydrogen storage and transport for use in hydrogen fuel cells can be dangerous.

Summary

Electrochemical cells generate electricity via redox reactions; E°cell = E°(cathode) − E°(anode).
Three main types: non-rechargeable (single-use), rechargeable (reversible), and fuel cells (continuous).
Lithium-ion cells are common rechargeable batteries with defined electrode reactions.
Hydrogen-oxygen fuel cells are efficient and produce only water, but storing hydrogen is difficult.
Electrochemical cells have advantages (portable energy, low emissions at point of use) and risks (waste, cost, safety).