R3.2.7 Secondary (rechargeable) cells

Reversibility of Redox Reactions in Electrochemical Cells

Understanding Reversibility in Redox Reactions

1. In chemistry, a redox reaction involves the transfer of electrons between two species.
2. One species is oxidized (loses electrons), while the other is reduced (gains electrons).
3. In some cases, these reactions are reversible, meaning the products can be converted back into reactants by applying an external energy source, such as an electric current.

How Reversibility Works

1. Reversibility in redox reactions is the basis of secondary (rechargeable) electrochemical cells.
2. During discharge, the cell converts chemical energy into electrical energy through spontaneous redox reactions.
3. When an external voltage is applied during charging, the redox reactions are reversed, restoring the reactants.
   • Discharge: Spontaneous redox reactions produce electrical energy.
   • Charge: Non-spontaneous redox reactions occur, driven by external electrical energy.

Examples of Reversible Redox Systems

1. Lead-Acid Batteries

• Lead-acid batteries are commonly used in cars and backup power systems.
• These batteries consist of a lead anode (negative electrode) and a lead(IV) oxide cathode (positive electrode) immersed in sulfuric acid.

Discharge (Powering the Device)

During discharge, the following reactions occur:

• Anode (Oxidation): $$\text{Pb(s)}+{\text{HSO}}_4^{-}(aq)\to {\text{PbSO}}_4(s)+{\text{H}}^{+}(aq)+2e^{-}$$
• Cathode (Reduction): $${\text{PbO}}_2(s)+3{\text{H}}^{+}(aq)+{\text{HSO}}_4^{-}(aq)+2e^{-}\to {\text{PbSO}}_4(s)+2{\text{H}}_2\text{O}(l)$$
• Overall Cell Reaction: $$\text{Pb(s)}+{\text{PbO}}_2(s)+2{\text{H}}_2{\text{SO}}_4(aq)\to 2{\text{PbSO}}_4(s)+2{\text{H}}_2\text{O}(l)$$

Charge (Recharging the Battery)

When an external voltage is applied, the reactions are reversed:

• Anode (Reduction): $${\text{PbSO}}_4(s)+{\text{H}}^{+}(aq)+2e^{-}\to \text{Pb(s)}+{\text{HSO}}_4^{-}(aq)$$
• Cathode (Oxidation): $${\text{PbSO}}_4(s)+2{\text{H}}_2\text{O}(l)\to {\text{PbO}}_2(s)+3{\text{H}}^{+}(aq)+{\text{HSO}}_4^{-}(aq)+2e^{-}$$
• Overall Cell Reaction: $$2{\text{PbSO}}_4(s)+2{\text{H}}_2\text{O}(l)\to \text{Pb(s)}+{\text{PbO}}_2(s)+2{\text{H}}_2{\text{SO}}_4(aq)$$

2. Nickel-Cadmium (NiCd) Cells

• Nickel-cadmium cells are another type of rechargeable battery, often used in portable electronic devices and power tools.
• These batteries consist of a cadmium anode and a nickel(III) oxide-hydroxide cathode in an alkaline electrolyte (usually potassium hydroxide).

Discharge (Powering the Device)

During discharge, the reactions are:

• Anode (Oxidation): $$\text{Cd(s)}+2{\text{OH}}^{-}(aq)\to {\text{Cd(OH)}}_2(s)+2e^{-}$$
• Cathode (Reduction): $$\text{NiO(OH)}(s)+{\text{H}}_2\text{O}(l)+e^{-}\to {\text{Ni(OH)}}_2(s)+{\text{OH}}^{-}(aq)$$
• Overall Cell Reaction: $$\text{Cd(s)}+2\text{NiO(OH)}(s)+2{\text{H}}_2\text{O}(l)\to {\text{Cd(OH)}}_2(s)+2{\text{Ni(OH)}}_2(s)$$

Charge (Recharging the Battery)

When an external voltage is applied, the reactions are reversed:

• Anode (Reduction): $${\text{Cd(OH)}}_2(s)+2e^{-}\to \text{Cd(s)}+2{\text{OH}}^{-}(aq)$$
• Cathode (Oxidation): $${\text{Ni(OH)}}_2(s)+{\text{OH}}^{-}(aq)\to \text{NiO(OH)}(s)+{\text{H}}_2\text{O}(l)+e^{-}$$
• Overall Cell Reaction: $${\text{Cd(OH)}}_2(s)+2{\text{Ni(OH)}}_2(s)\to \text{Cd(s)}+2\text{NiO(OH)}(s)+2{\text{H}}_2\text{O}(l)$$

3. Lithium-Ion Batteries

• Lithium-ion batteries are commonly used in portable electronics, electric vehicles, and energy storage systems due to their high energy density and rechargeability.
• These batteries consist of a lithium cobalt oxide (LiCoO₂) cathode and a graphite anode, with a lithium salt electrolyte.

Discharge (Powering the Device)

During discharge, lithium ions move from the anode to the cathode, generating electrical energy:

• Anode (Oxidation):
  $${\text{LiC}}_6\to C_6+{\text{Li}}^{+}+e^{-}$$
• Cathode (Reduction):
  $${\text{LiCoO}}_2+{\text{Li}}^{+}+e^{-}\to {\text{Li}}_2{\text{CoO}}_2$$
• Overall Cell Reaction:
  $${\text{LiC}}_6+{\text{LiCoO}}_2\to C_6+{\text{Li}}_2{\text{CoO}}_2$$

Charge (Recharging the Battery)

During charging, an external power source drives the reverse reactions, storing energy by moving lithium ions back to the anode:

• Anode (Reduction):
  $$C_6+{\text{Li}}^{+}+e^{-}\to {\text{LiC}}_6$$
• Cathode (Oxidation):
  $${\text{Li}}_2{\text{CoO}}_2\to {\text{LiCoO}}_2+{\text{Li}}^{+}+e^{-}$$
• Overall Cell Reaction:
  $$C_6+{\text{Li}}_2{\text{CoO}}_2\to {\text{LiC}}_6+{\text{LiCoO}}_2$$

Advantages and Disadvantages of Fuel Cells, Primary Cells, and Secondary Cells

Each type serves specific purposes, with fuel cells excelling in efficiency, primary cells in convenience, and secondary cells in sustainability.

Reflection