Wiki Category: science

How do Lead Acid Batteries Work?


The lead acid battery was the first rechargeable battery, first created by French physicist Gaston Plante in 1859. Despite having been outclassed by lithium ion in terms of energy density, still remains widely used for its low cost and ability to supply large currents. Lead-acid batteries are the second most manufactured battery type after lithium ion.


A fully charged lead acid battery consists of a cathode made of lead dioxide, an anode of lead metal, and sulfuric acid as the electrolyte. The lead metal plate being the anode we know that the lead must be undergoing oxidation.


The Pb2+ then combines with HSO4 in solution to form a solid precipitate of PbSO4 around the lead anode. This releases the hydrogen attached to the sulfate which would make the full half reaction occurring at the anode:

    \[Pb+HSO_4^-\rightarrow PbSO^4+H^++2e^-\]

At the lead dioxide cathode we know the lead dioxide must be getting reduced. The lead atom in lead dioxide has an oxidation number of +4 due to its two oxygens which means that it can be reduced by also becoming Pb2+. To facilitate this, the hydrogen ions produced by the reaction in the anode are consumed in this reaction to remove the two oxygens as two waters.


This would make the overall reaction:


As you can see, in this reaction the acid is consumed and converted to water which means, in its discharged state, a lead acid battery would have both of its electrodes coated in lead sulfate and the sulfuric acid would have been almost entirely consumed.

Diagram displaying the electrodes in a simple lead acid battery cell

When the battery is charged the reaction is reversed and the two electrodes return to their original state. Of course, this process is not perfect which leads the battery losing some capacity every charge cycle. One mechanism by which this occurs is through sulfation in which the lead sulfate produced by the reaction forms crystals in such a way that renders the lead sulfate unable to participate in the reaction again. In some cases, sulfation can be reversed, but not always.


While early lead acid batteries were as simple as two plates submerged in acid, modern lead acid batteries use a variety of manufacturing techniques to maximize power and minimize electrode corrosion.


One feature of every modern lead acid battery cell is that each electrode being a single plate, such as one you would use in a simple lab set up, each electrode consists of several attached plates. This increases the surface area of the battery which allows more mols of the reaction to occur simultaneously which increases the flow of electrons across the battery. Note that this does not increase the voltage of the cell, the voltage will always be more or less the same despite configuration. But it does allow more electrons to travel at a given time. The increase in surface area could thus be thought of as decreasing the internal resistance of the battery which allows for a higher maximum current. The standard potential of one cell is 2.05V. Thus, the average 12V lead acid battery would be made up of 6 of these cells.

Diagram displaying the components of a lead-acid battery

In the above image you can see how each cell consists of interlocking plates that are kept apart by a separator that allows ions to pass through. Note that each plate is only attached to plates of the same electrode.


The anode needs to be made of lead, however, pure lead is too soft to maintain its shape. Therefore, lead alloys are used to strengthen the plates and improve the electrical properties of the electrodes. Some added metals include tin, selenium, calcium, and antimony. Each metal contributes differently to the performance of the battery and thus batteries with different intended purposes will use different alloys. For example,


The cathode is made up of plates coated with a lead oxide paste. Modern lead acid batteries include various additives to the paste to improve performance. One example is barium sulfate which acts as a seed crystal for the formation of lead sulfate crystals. When the lead sulfate forms in these crystals, it is easier for the reverse reaction to occur and convert the lead sulfate back to lead oxide. Other additives can be introduced to the electrolyte to boost performance.


Another innovation in the lead acid battery was improvement of the separator. Many lead-acid batteries now use absorbent glass mat separators(AGMs). These are mats made of glass fibers which soak up the electrolyte. This is advantageous compared to a flooded lead acid battery because if the battery is punctured, no acid will leak out. Also, since the electrolyte is confined to the glass mat, the battery can now experience more turbulence without potentially damaging the battery.


Most of the world’s lead-acid batteries are used to start the engines and power the electrical systems of automobiles. Most of the stored charge of the lead acid battery isn’t directly used as, once the engine has started, the vehicle’s alternator converts some of the engines power into electricity which powers the vehicle’s electrical systems and charges the battery.
Lead-acid batteries can also be used to power small, fully electric vehicles such as golf carts, electric scooters, electric bikes, and electric wheelchairs.
Lead-acid batteries are also used to store backup power for telecommunication systems.

References/Further Reading

Lead-Acid Batteries Tutorial

Battery University: How does the lead-acid battery work?
Lead Acid Battery Chart

Battery Charging and Maintenance (images)

How do Batteries Work?


Broadly speaking, a battery is anything that can store energy and then release that energy in a controlled manner. Our bodies can be considered batteries in the way they store chemical energy in the form of fat and other substances. Springs in mechanical systems are like batteries that store and release mechanical energy. Usually however, when we say battery, we are referring to electrochemical batteries, batteries that store electric energy as chemical potential energy. These batteries are also referred to as electrochemical cells. These batteries power nearly every electronic device that is not always connected to the electrical grid. From flashlights to tablets to cars, all these things need electrochemical batteries to function.

So how does a battery turn chemical energy into electrical energy and vice versa? Well at first, batteries could only discharge and unlike your phone battery could not be recharged, so lets look at how the first electrochemical cells functioned.

The Voltaic Pile

The first electrochemical cell was created by Italian physicist Alessandro Volta which is now known as the voltaic pile. The voltaic pile was a stack of individual cells, the smallest functioning unit of a battery, each of which consisted of a zinc and copper plate separated by a salt solution-soaked cloth. Connected to the top and bottom of the stack is the wire that carries the electric current.

In each individual cell can be seen the three fundamental components in every battery: The anode, the cathode and the electrolyte. In this pile, copper is acting as the anode and zinc is acting as the cathode while the saltwater in the paper is the electrolyte. Electricity can be understood as the flow of electrons through conductive materials like wires. The flow of electrons is caused by a difference in electric potential energy. Copper is more electronegative than zinc which means it attracts electrons to itself more strongly. The resulting potential difference causes zinc to undergo a spontaneous half-reaction known as reduction. A reduction reaction is a reaction where electrons are removed from a chemical species.

    \[Zn\rightarrow{\rm Zn}^{2+}+2e^-\]

This is known as a half reaction because another reaction is usually required to consume the electrons since electrons are more stable when they exist in atoms. A reaction that adds electrons to a chemical species is called an oxidation reaction. The oxidation and reduction reaction occur at the cathode and anode respectively, with the reduction reaction producing electrons and flow from the anode and through the circuit to be consumed in the oxidation reaction occurring at the cathode. The oxidation half reaction occurring at the cathode in the voltaic pile is shown below.

    \[2H_2O+2e^-\rightarrow2{\rm OH}^-+H^2\]

In this reaction, the excess electrons on the copper plate attack water to turn it into hydroxide and hydrogen gas. The element being oxidized in this oxidation reaction is oxygen since it gains an electron. These two half reactions can be added together to form one complete reaction.

    \[Zn\rightarrow{\rm Zn}^{2+}+2e^-\]

    \[\underline{+\ 2H_2O+2e^-\rightarrow2{\rm OH}^-+H^2}\]


This complete reaction is known as a redox(reduction + oxidation) reaction. This reaction continues as long as there is zinc to be dissolved and a path allowing the flow of electrons. Note that this reaction does not include copper but the copper is necessary as the difference in electronegativity between copper and zinc acts as a driving force in this reaction. After this example, another will be shown of an electrolytic cell where copper is involved in the reaction.

Diagram displaying reactions–

The Galvanic Cell

In high school and college chemistry classes, one often required lab is the construction of a galvanic cell, similar to what is shown below.

In this cell we can see that copper, unlike in the voltaic pile example, is acting as the anode while silver is acting as the cathode. Also both electrodes exist in their own solution containing their respective ions to act as the electrolyte. The electrolyte in this cell will be participating in the redox reaction. Each half reaction and the full reaction are below.

    \[Cu\rightarrow{\rm Cu}^{2+}+2e^-\]

    \[\underline{2{\rm Ag}^++2e^-\rightarrow2Ag}\]

    \[Cu+2{\rm Ag}^+\rightarrow{\rm Cu}^{2+}+2Ag\]

We can see that, copper is the anode in this cell because it is being oxidized(gaining electrons) in the reaction and vice versa for silver.

Because the electrolytes are now directly involved in the chemical reaction, they must be separated into half cells so that the reaction between Cu and Ag+ doesn’t occur at the surface of the copper electrode. It is important that this doesn’t occur because then the electron transfer would be direct and not force the electrons to travel around the wire.

Since there are now two solutions unlike in the voltaic pile, a salt bridge is necessary to allow the balance of charge. This salt bridge allows spectator ions(ions not involved in any reaction) to travel between the solutions to balance the charge. In this case the spectator ions are Na+ and No3-.

To summarize, batteries produce electrical power by separating the reactants of a redox reaction such that it can only occur if the electron transfer needed can only occur by traveling through an electric circuit. Which chemicals used and how they are separated vary across the different battery types, each setup carrying with it different pros and cons.

Sources/Further Reading