Lead–acid battery

Lead-acid batteries, invented in 1859 by French physicist Gaston Planté, are a type of galvanic cell and are the most commonly used rechargeable batteries today. They also represent the oldest design with one of the worst energy-to-weight ratios, although the power-to-weight ratio can be quite good. Also, the energy-to-volume ratio is good compared to other types of batteries. They are cheap and can supply high surge currents needed in starter motors. Every reasonably modern car uses a lead-acid battery for this purpose. They are also used in vehicles such as forklifts, in which the low energy-to-weight ratio may in fact be considered a benefit since the battery can be used as a counterweight.

Lead-acid car batteries consist of six cells of 2 V nominal voltage. Each cell contains (in the charged state) electrodes of lead metal (Pb) and lead (IV) oxide (PbO₂) in an electrolyte of about 37 % w/w sulfuric acid (H₂SO₄). Modern designs have gelified electrolytes. In the discharged state both electrodes turn into lead(II) sulfate and the electrolyte turns into water. (This is why discharged lead-acid batteries can freeze.)

Lead acid batteries for automotive use are not designed for deep discharge and should always be kept at maximum charge, using constant voltage at 13.8 V (for six element car batteries). Their capacity will severely suffer from deep cycling. Specially designed deep-cycle cells are much less susceptible to this problem, and are required for applications where the batteries are regularly discharged.

Quiescent(open-circuit) voltage at full battery: 12.6 V
Unloading-end voltage: 11.8 V
Charge with 13.2-14.4 V
Gassing voltage: 14.4 V
Continuous-preservation charge with max. 13.2 V
After full charge the terminal voltage will drop quickly to 13.2 V and then slowly to 12.6 V.

The energy to weight ratio, or specific energy, is in the range of 108 kJ/kg (30 Wh/kg).

The chemical reactions are (charged to discharged):

Anode (oxidation): ${\mbox{Pb}}(s)+{\mbox{SO}}_{4}^{2-}(aq)\leftrightarrow {\mbox{PbSO}}_{4}(s)+2e^{-}\quad \epsilon ^{o}=0.356V$

Cathode (reduction): ${\mbox{PbO}}_{2}(s)+{\mbox{SO}}_{4}^{2-}(aq)+4{\mbox{H}}^{+}+2e^{-}\leftrightarrow {\mbox{PbSO}}_{4}(s)+2{\mbox{H}}_{2}{\mbox{O}}(l)\quad \epsilon ^{o}=1.685V$

Because of the open cells with liquid electrolyte in most cheap car batteries, overcharging with excessive charging voltages will generate oxygen and hydrogen gas, forming an extremely explosive mix. This should be avoided. Caution must also be observed because of the extremely corrosive nature of sulphuric acid.

Automotive and other applications

A chemical compound in the form of tablets can be added to each cell to reduce sulfate build up, and improve battery condition, however the effectiveness of such treatments is subject to debate.

See Car battery

Wet cells designed for deep discharge are commonly used in golf carts and other battery electric vehicles, large backup power supplies for telephone and computer centers and off-grid household electric power systems.

Absorbed glass mat (AGM) cells are also used in battery electric vehicles.

Gel cells are used in back-up power supplies for alarm and smaller computer systems, and for electric scooters and electrified bicycles.

References

U.S. Department of Energy, Primer On Lead-Acid Storage Batteries [1] (pdf).

Automotive and other applications

See also

References