Lead Storage Batteries

Lead storage batteries, first created by Gaston Plante in 1859, were the first rechargeable battery created. Despite being a very old technology, they are still quite common because they are relatively inexpensive and can provide a large current (lots of electrons flowing at the same time). Their most common use today is the battery found in your gasoline burning car. (Electric cars use lithium-ion batteries.)

The Structure

Lead storage batteries take advantage of the fact that lead has three common charges: 0, +2 and +4, with the middle charge being more stable than the higher, +4 charge. Each cell in the battery is composed of two grids, or thin boxes made of conducting material. The first grid is filled with powdered lead, the other with lead oxide and the two are then placed in a bath of aqueous sulfuric acid ($H_{2}S{O}_4$).

Because sulfuric acid is diprotic and the first acid is strong, the solution consists primarily of $HS{O}_4^{-1}$ ions and $H^{+1}$ ions.

The battery is comprised of 6 of these cells, that is 12 grids alternating between $Pb$ and $Pb{O}_2$.

http://www.vehiclemaintenanceandrepairs.com/car-battery-facts-101/

The Reactions (simplified)

In simplest terms, the reactions in a lead storage battery are simply an exchange of electrons between the +4 lead and the neutral lead.

$P{b}^{+4}+2e^{-1}\to P{b}^{+2}$

$Pb\to P{b}^{+2}+2e^{-1}$

These two half-reactions give 2 V, so the 6 cells together provide 12V.

Recharging (simplified)

Recharging the battery requires us to force the two reactions above to occur backwards:

$P{b}^{+2}\to P{b}^{+4}+2e^{-1}$

$P{b}^{+2}+2e^{-1}\to Pb$

We do this by applying MORE than 12V in the “wrong” direction. In your car, this is done by the alternator, which is essentially an electric power plant that runs off of the engine.

In fact, the alternator is what provides the electricity to all of the parts of your car (A/C, radio, lights, etc) when the engine is running. Only when the engine is off, do the electrical systems of your car pull electricity from the battery. That’s why your fuel efficiency is lower in the summer when you run the A/C than it is in the winter. The engine has to burn extra gas to make the electricity to run the A/C.

The Problem with Recharging

When we recharge the battery, we force the $P{b}^{+2}$ ions to either take two electrons and become $Pb$ or to give up two electrons and become $P{b}^{+4}$. However, we have a problem. Lead II ions ($P{b}^{+2}$) are somewhat soluble in water. That means that these ions could dissolve into the electrolyte solution in the battery. If they did that, they would leave their positions in the grids. Then, when we try to force those reactions to go backwards, those ions will not be there to react.

We needed to find a way to make sure that the ions were still locked in the grids when the alternator kicked in. The solution was to make the $P{b}^{+2}$ ions insoluble. This is done by having sulfate ($S{O}_4^{-2}$) ions present in the solution. Lead II sulfate ($PbS{O}_4$) is quite insoluble, so when the $P{b}^{+2}$ ions are formed, they react with the sulfate ions, make $PbS{O}_4$ and stay put, waiting to be “recharged.”

The Reactions - Complete

Once we understood the need to make the $P{b}^{+2}$ ions insoluble, we addressed this need by using sulfuric acid. The resulting (complete) reactions are these:

$P{b}_{(s)}+HS{O}_{4(aq)}^{-1}\to PbS{O}_{4(s)}+H_{(aq)}^{+1}+2e^{-1}$

$Pb{O}_{2(s)}+HS{O}_{4(aq)}^{-1}+3H_{(aq)}^{+1}+2e^{-1}\to PbS{O}_{4(s)}+2H_2{O}_{(l)}$