Lead-Acid Batteries Chemistry Tutorial

Key Concepts

A lead-acid battery is made up of a number of lead-acid galvanic (voltaic) cells connected up in series.

When a lead-acid cell is producing electricity (discharging) it is converting chemical energy into electrical energy.

Discharging a lead-acid battery is a spontaneous redox reaction.

When a single lead-acid galvanic cell is discharging, it produces about 2 volts.

6 lead-acid galvanic cells in series produce 12 volts.
The battery in a petrol or diesel car is a 12 volt lead-acid battery.

Lead-acid cells are rechargeable because the reaction products do not leave the electrodes.

A lead-acid galvanic cell can be recharged by connecting the :
(i) negative terminal of a battery charger to the negative terminal of the galvanic cell
(ii) positive terminal of a battery charger to the positive terminal of the galvanic cell.

(iii) A lead-acid battery can therefore be recharged in a similar way.

Recharging a lead-acid galvanic cell (or a battery) converts electrical energy from the recharger back into chemical energy stored in the galvanic cell (or battery).

Recharging a lead-acid cell is therfore a non-spontaneous electrolytic process.

Galvanic cells which can be recharged are known as secondary cells.
Lead-acid cells are secondary cells.

Please do not block ads on this website.
No ads = no money for us = no free stuff for you!

A Lead-Acid Galvanic Cell Producing Electricity (discharging)

Discharge of a lead-acid cell:

is a spontaneous redox reaction (E_(redox) is positive)

converts chemical energy stored in the lead, lead dioxide and sulfuric acid into electrical energy

Below is a diagram of a lead-acid cell during discharge:

anode (-)

cathode (+)

e^- →

electrolyte

At the anode during discharge:

Electrode: spongy lead, Pb_(s) in a lead frame (porous lead plate)
Oxidation occurs ar the anode.

Lead is oxidized at the anode:
Pb_(s) → Pb²⁺_(aq) + 2e^-

Pb²⁺ reacts with SO₄^2- in sulfuric acid to produce PbSO_4(s) as a precipitate:
Pb²⁺_(aq) + SO₄^2-_(aq) → PbSO_4(s)

PbSO_4(s) sticks to the lead frame of the anode.
The overall reaction at the anode is
Pb_(s)+ SO₄^2-_(aq) → PbSO_4(s)+ 2e^-

Electrons are produced at the anode.

Anode is negative.

Electrons migrate from anode (-) to cathode (+) along wires connecting the electrodes.

At the cathode during discharge:

Electrode: lead frame covered with lead dioxide, PbO_2(s)

Reduction occurs at the cathode.

Lead in lead dioxide is reduced at the cathode by protons from the sulfuric acid:
PbO_2(s) + 4H⁺_(aq) + 2e^- → Pb²⁺_(aq) + 2H₂O_(l)

Pb²⁺ reacts with SO₄^2- in sulfuric acid to produce PbSO_4(s) as a precipitate:
Pb²⁺_(aq) + SO₄^2-_(aq) → PbSO_4(s)

PbSO_4(s) sticks to the lead frame of the cathode.

The overall reaction at the cathode is
PbO_2(s) + 4H⁺_(aq) + 2e^- + SO₄^2-_(aq) → PbSO_4(s) + 2H₂O_(l)

Electrons are consumed at the cathode.

Cathode is positive.

Protons, H⁺_(aq), are consumed at the cathode.

Wen can use the oxidation reaction at the anode and the reduction reaction that occurs at the cathode to write an overall redox equation for the lead-acid cell:

anode reaction:	Pb_(s) + SO₄^2-_(aq)	→	PbSO_4(s)+ 2e^-
cathode reaction:	PbO_2(s) + 4H⁺_(aq) + SO₄^2-_(aq) + 2e^-	→	PbSO_4(s) + 2H₂O_(l)

overall reaction:	Pb_(s) + PbO_2(s) + 4H⁺_(aq) + 2SO₄^2-_(aq)	→	2PbSO_4(s) + 2H₂O_(l)	E ≈ +2 V

As the lead-acid cell discharges:

PbSO₄ precipitates out and deposits on both the anode and the cathode.

H⁺ from the electrolyte (H₂SO_4(aq)) is being used to produce water at the cathode.

Concentration of H⁺ will be decreased over time (concentration of H₂SO_(aq) decreases).

pH of the electrolyte (H₂SO_4(aq)) will increase.

Connecting lead-acid galvanic cells in a series to make lead-acid batteries allows us to make batteries with voltages greater than 2 V:

number of cells in series	2	3	4	5	6
approximate voltage produced	2 × 2 = 4 V	3 × 2 = 6 V	4 × 2 = 8 V	5 × 2 = 10 V	6 × 2 = 12 V

Do you know this?

Join AUS-e-TUTE!

Play the game now!

Recharging a Lead-Acid Galvanic Cell

Recharging a lead-acid cell:

is a non-spontaneous redox reaction (E_(redox) is negative), that is, an electrolytic process

requires an input of slightly more than 2 volts per cell to drive the spontaneous reactions in the reverse direction

converts electrical energy back into chemical energy which is stored in the lead, lead dioxide and sulfuric acid in the cell

The diagram below shows a lead-acid cell during recharge:

cathode (-)

anode (+)

← e^-

electrolyte

At the cathode during recharge:

electrons are being pushed into the cathode from the recharger

cathode is negative

reduction occurs at the cathode

lead in lead sulfate sticking to the electrode is reduced back to Pb_(s):
PbSO_4(s)+ 2e^- → Pb_(s)+ SO₄^2-_(aq)

At the anode during recharge:

electrons are being pulled out of the anode by the recharger

anode is positive

oxidation occurs at the anode

lead in lead sulfate sticking to the electrode is oxidized by water to reform lead dioxide:
PbSO_4(s) + 2H₂O_(l) → PbO_2(s) + 4H⁺_(aq) + SO₄^2-_(aq) + 2e^-

protons, H⁺, are produced at the anode

electrons migrate from anode (+) to cathode (-)

Using the half-equations for the reactions occuring at the anode and cathode we can write an overall redox reaction for the lead-acid cell during recharge:

cathode reaction:	PbSO_4(s)+ 2e^-	→	Pb_(s) + SO₄^2-_(aq)
anode reaction:	PbSO_4(s) + 2H₂O_(l)	→	PbO_2(s) + 4H⁺_(aq) + SO₄^2-_(aq) + 2e^-

overall reaction:	2PbSO_4(s) + 2H₂O_(l)	→	Pb_(s) + PbO_2(s) + 4H⁺_(aq) + 2SO₄^2-_(aq)	E ≈ -2 V

As the lead-acid cell recharges:

PbSO_4(s) on each electrode is removed.

Concentration of H⁺ increases.

pH of the electrolyte (H₂SO_4(aq)) decreases.

Recharging a lead-acid battery requires a minimum of 2 V per cell:

number of cells in series	2	3	4	5	6
recharge voltage required	> 4 V	> 6 V	> 8 V	> 10 V	> 12 V

Overcharging a battery electrolyzes water, producing hydrogen gas and oxygen gas.

The bubbles of gas degrade the surfaces of the electrodes causing the PbSO_4(s) to fall off the electrodes.

This reduces the capacity of the cell.

Do you understand this?

Join AUS-e-TUTE!

Take the test now!