Batteries

Batteries

Batteries are indispensable in the modern electronic world. For example, Li – ion batteries are used in cell phones, dry cell in flashlight etc…. These batteries are used as a source of direct current at a constant voltage. We can classify them into primary batteries (non – rechargeable) and secondary batteries (rechargeable). In this section, we will briefly discuss the electrochemistry of some batteries.

Leclanche cell

Anode : Zinc container

Cathode : Graphite rod in contact with MnO₂

Electrolyte : ammonium chloride and zinc chloride in water

Emf of the cell is about 1.5V

Cell reaction

Oxidation at anode

Zn (s) → Zn²⁺ (aq)+2e^-                     .....(1)

Reduction at cathode

2 NH₄⁺ (aq) + 2e^- → 2NH₃(aq) + H₂ (g)                 .....(2)

The hydrogen gas is oxidised to water by MnO₂

H₂ (g) + 2 MnO₂ (s) → Mn₂ O₃ (s) + H₂O (l)                   .....(3)

Equation (1) + (2) + (3) gives the overall redox reaction

Zn (s) + 2NH₄⁺ (aq) + 2 MnO₂ (s) → Zn²⁺ (aq) + Mn₂ O₃(s) + H₂O (l) + 2NH₃

...... (4)

Ammonia produced at the cathode combines with Zn²⁺ to form a complex ion [Zn (NH₃ )₄ ]²⁺ (aq) . As the reaction proceeds, the concentration of NH₄⁺ will decrease and the aqueous NH₃ will increase which lead to the decrease in the emf of cell.

Mercury button cell

Anode : zinc amalgamated with mercury

Cathode : HgO mixed with graphite

Electrolyte : Paste of KOH and ZnO

Oxidation occurs at anode : 

Reduction occurs at cathode : 

Overall reaction : Zn (s) + HgO (s) → ZnO (s) + Hg (l)

Cell emf : about 1.35V.

Uses : It has higher capacity and longer life. Used in pacemakers, electronic watches, cameras etc…

Secondary batteries

We have already learnt that the electrochemical reactions which take place in a galvanic cell may be reversed by applying a potential slightly greater than the emf generated by the cell. This principle is used in secondary batteries to regenerate the original reactants. Let us understand the function of secondary cell by considering the lead storage battery as an example

Lead storage battery

Anode : spongy lead

Cathode : lead plate bearing PbO₂

Electrolyte : 38% by mass of H₂ SO₄ with density 1.2g / mL.

Oxidation occurs at the anode

Pb(s) → Pb²⁺ (aq)+2e^-              .....(1)

The Pb²⁺ ions combine with SO₄^-2 to from PbSO₄ precipitate.

Pb²⁺ (aq) + SO₄ ^2−(aq) → PbSO₄ (s)                   .....(2)

Reduction occurs at the cathode

PbO₂ (s) + 4 H⁺ (aq) + 2e^- → Pb²⁺ (aq) + 2H₂O (l)                 ......(3)

The Pb²⁺ ions also combine with SO₄^2- ions from sulphuric acid to form precipitate.

PbSO₄ Pb²⁺ (aq) + SO₄ ^2− (aq) → PbSO₄ (s)                  .....(4)

The Overall reactions is

Equation (1) + (2) + (3) + (4)

Pb (s) + PbO₂ (s) + 4H ⁺ (aq) + 2SO₄ ^2− (aq) → 2 PbSO₄ (s) + 2H₂O (l )

The emf of a single cell is about 2V . Usually six such cells are combined in series to produce 12volt

The emf of the cell depends on the concentration of H₂ SO₄ . As the cell reaction uses SO₄^2− ions, the concentration H₂ SO₄ decreases. When the cell potential falls to about 1.8V, the cell has to be recharged.

Recharge of the cell

As said earlier, a potential greater than 2V is applied across the electrodes, the cell reactions that take place during the discharge process are reversed. During recharge process, the role of anode and cathode is reversed and H₂ SO₄ is regenerated.

Oxidation occurs at the cathode ( now act as anode)

Reduction occurs at the anode (now act as cathode) PbSO₄ (s) + 2e^- → Pb(s) + SO₄^2-(aq)

Overall reaction

2PbSO₄(s) + 2H₂O (l ) → Pb (s) + PbO₂ (s) + 4H⁺ (aq) + 2SO₄^2- (aq).

Thus, the overall cell reaction is exactly the reverse of the redox reaction which takes place while discharging .

Uses:

Used in automobiles, trains, inverters etc…

The lithium – ion Battery

Anode : Porus graphite

Cathode : transition metal oxide such as CoO₂.

Electrolyte : Lithium salt in an organic solvent

At the anode oxidation occurs

Li (s) → Li⁺ (aq) + e^-

At the cathode reduction occurs

Li⁺ + CoO₂ (s) + e^- → Li CoO₂ (s)

Overall reactions

Li (s) + CoO₂ → LiCoO₂ (s)

Both electrodes allow Li⁺ ions to move in and out of their structures.

During discharge, the Li⁺ ions produced at the anode move towards cathode through the non – aqueous electrolyte. When a potential greater than the emf produced by the cell is applied across the electrode, the cell reaction is reversed and now the Li⁺ ions move from cathode to anode where they become embedded on the porous graphite electrode. This is known as intercalation.

Uses :

Used in cellular phones, laptops, computers, digital cameras, etc…

Fuel cell

The galvanic cell in which the energy of combustion of fuels is directly converted into electrical energy is called the fuel cell. It requires a continuous supply of reactant to keep functioning. The general representation of a fuel cell is follows

Fuel | Electrode | Electrolyte | Electrode | Oxidant

Let us understand the function of fuel cell by considering hydrogen – oxygen fuel cell. In this case, hydrogen act as a fuel and oxygen as an oxidant and the electrolyte is aqueous KOH maintained at 200^oC and 20 – 40 atm. Porous graphite electrode containing Ni and NiO serves as the inert electrodes.

Hydrogen and oxygen gases are bubbled through the anode and cathode, respectively.

Oxidation occurs at the anode:

2H₂ (g ) + 4 OH^− (aq) → 4 H₂O (l) + 4 e^−

Reduction occurs at the cathode O ₂ (g) + 2 H₂O (l) + 4e^− → 4 OH^− (aq)

The overall reaction is 2H₂ (g) + O₂ (g) → 2H₂O (l)

The above reaction is the same as the hydrogen combustion reaction, however, they do not react directly ie., the oxidation and reduction reactions take place separately at the anode and cathode respectively like H₂ - O₂ fuel cell. Other fuel cells like propane –O₂ and methane O₂ have also been developed.