Lead–acid battery

The lead–acid battery is a type of rechargeable battery. First invented in 1859 by French physicist Gaston Planté, it was the first type of rechargeable battery ever created. Compared to the more modern rechargeable batteries, lead–acid batteries have relatively low energy density and heavier weight. Despite this, they are able to supply high surge currents. These features, along with their low cost, make them useful for motor vehicles in order to provide the high current required by starter motors. Lead–acid batteries suffer from relatively short cycle lifespan and overall lifespan, as well as long charging times; an average automotive battery takes anywhere between 6 to 12 hours to fully charge from a discharged state.
As they are not as expensive when compared to newer technologies, lead–acid batteries are widely used even when surge current is not important and other designs could provide higher energy densities. In 1999, lead–acid battery sales accounted for 40–50% of the value from batteries sold worldwide, equivalent to a manufacturing market value of about US$15 billion. Large-format lead–acid designs are widely used for storage in backup power supplies in telecommunications networks such as for cell sites, high-availability emergency power systems as used in hospitals, and stand-alone power systems. For these roles, modified versions of the standard cell may be used to improve storage times and reduce maintenance requirements. Gel cell and absorbed glass mat batteries are common in these roles, collectively known as valve-regulated lead–acid ''batteries''.
When charged, the battery's chemical energy is stored in the potential difference between metallic lead at the negative side and lead dioxide on the positive side.

History

The French scientist Nicolas Gautherot observed in 1801 that wires that had been used for electrolysis experiments would themselves provide a small amount of secondary current after the main battery had been disconnected. In 1859, Gaston Planté's lead–acid battery was the first battery that could be recharged by passing a reverse current through it. Planté's first model consisted of two lead sheets separated by rubber strips and rolled into a spiral and immersed in a solution containing about 10 percent sulfuric acid. His batteries were first used to power the lights in train carriages while stopped at a station. In 1881, Camille Alphonse Faure invented an improved version that consisted of a lead grid lattice into which a lead oxide paste was pressed, forming a plate. This design was easier to mass-produce. An early manufacturer of lead–acid batteries was Henri Tudor.
Using a gel electrolyte instead of a liquid allows the battery to be used in different positions without leaking. Gel electrolyte batteries for any position were first used in the late 1920s, and in the 1930s, portable suitcase radio sets allowed the cell to be mounted vertically or horizontally due to valve design. In the 1970s, the valve-regulated lead–acid, or sealed, battery was developed, including modern absorbed glass mat types, allowing operation in any position.
It was discovered early in 2011 that lead–acid batteries do in fact use some aspects of relativity to function, and to a lesser degree, liquid metal and molten-salt batteries such as the Ca-Sb and Sn-Bi also use this effect.

Electrochemistry

Discharge

In the discharged state, both the positive and negative plates become lead sulfate, and the electrolyte loses much of its dissolved sulfuric acid and becomes primarily water.
;Negative plate reaction:Pb + → + + 2e⁻
The release of two conduction electrons gives the lead electrode a negative charge.
As electrons accumulate, they create an electric field that attracts hydrogen ions and repels sulfate ions, leading to a double layer near the surface. The hydrogen ions screen the charged electrode from the solution, which limits further reaction unless charge is allowed to flow out of the electrode.
;Positive plate reaction: + + 3 + 2e⁻ → + 2
taking advantage of the metallic conductivity of lead dioxide|.
;The total reaction can be written as
The net energy released per mole of Pb converted to is approximately 400 kJ, corresponding to the formation of 36 g of water. The sum of the molecular masses of the reactants is 642.6 g/mole, so theoretically a cell can produce two faradays of charge from 642.6 g of reactants, or 83.4 ampere-hours per kilogram for a 2-volt cell. This comes to 167 watt-hours per kilogram of reactants, but in practice, a lead–acid cell gives only 30–40 watt-hours per kilogram of battery, due to the mass of the water and other constituent part.
Another form for discharging reaction:
Negative plate:
Pb + H₂SO₄ → PbSO₄ + H₂ + 2e^-
Positive plate:
PbO₂ + H₂SO₄ + H₂ + 2e^- → PbSO₄ + 2H₂O

Charging

In the fully charged state, the negative plate consists of lead and the positive plate is lead dioxide. The electrolyte solution has a higher concentration of aqueous sulfuric acid, which stores most of the chemical energy.
Overcharging with high charging voltages generates oxygen and hydrogen gas by electrolysis of water, which bubbles out and is lost. The design of some types of lead–acid battery allows the electrolyte level to be inspected and topped up with pure water to replace any that has been lost this way.

Effect of charge level on freezing point

Because of freezing-point depression, the electrolyte is more likely to freeze in a cold environment when the battery has a low charge and a correspondingly low sulfuric acid concentration.

Ion motion

During discharge, produced at the negative plates moves into the electrolyte solution and is then consumed at the positive plates, while is consumed at both plates. The reverse occurs during the charge. This motion can be electrically-driven proton flow, or by diffusion through the medium, or by the flow of a liquid electrolyte medium. Since the electrolyte density is greater when the sulfuric acid concentration is higher, the liquid will tend to circulate by convection. Therefore, a liquid-medium cell tends to rapidly discharge and rapidly charge more efficiently than an otherwise-similar gel cell.

Measuring the charge level

Because the electrolyte takes part in the charge-discharge reaction, this battery has one major advantage over other chemistries: it is relatively simple to determine the state of charge by merely measuring the specific gravity of the electrolyte; the specific gravity falls as the battery discharges. Some battery designs include a simple hydrometer using colored floating balls of differing density. When used in diesel-electric submarines, the specific gravity was regularly measured and written on a blackboard in the control room to indicate how much longer the boat could remain submerged.
The battery's open-circuit voltage can also be used to gauge the state of charge. If the connections to the individual cells are accessible, then the state of charge of each cell can be determined which can provide a guide as to the state of health of the battery as a whole; otherwise, the overall battery voltage may be assessed.

Voltages for common usage

is a three-stage charging procedure for lead–acid batteries. A lead–acid battery's nominal voltage is 2.1 V for each cell. For a single cell, the voltage can range from 1.8 V loaded at full discharge, to 2.10 V in an open circuit at full charge.
Float voltage varies depending on battery type, and ranges from 1.8 V to 2.27 V. Equalization voltage, and charging voltage for sulfated cells, can range from 2.67 V to almost 3 V. Specific values for a given battery depend on the design and manufacturer recommendations, and are usually given at a baseline temperature of, requiring adjustment for ambient conditions. IEEE Standard 485-2020 is the industry's recommended practice for sizing lead–acid batteries in stationary applications.

Construction

Plates

The lead–acid cell can be demonstrated using sheet lead plates for the two electrodes. However, such a construction produces only around one ampere for roughly postcard-sized plates, and for only a few minutes.
Gaston Planté found a way to provide a much larger effective surface area. In Planté's design, the positive and negative plates were formed of two spirals of lead foil, separated with a sheet of cloth and coiled up. The cells initially had low capacity, so a slow process of forming was required to corrode the lead foils, creating lead dioxide on the plates and roughening them to increase surface area. Initially, this process used electricity from primary batteries; when generators became available after 1870, the cost of producing batteries greatly declined. Planté plates are still used in some stationary applications, where the plates are mechanically grooved to increase their surface area.
In 1880, Camille Alphonse Faure patented a method of coating a lead grid with a paste of lead oxides, sulfuric acid, and water, followed by curing phase in which the plates were exposed to gentle heat in a high-humidity environment. The curing process changed the paste into a mixture of lead sulfates which adhered to the lead plate. Then, during the battery's initial charge, the cured paste on the plates was converted into electrochemically active material. Faure's process significantly reduced the time and cost to manufacture lead–acid batteries, and gave a substantial increase in capacity compared with Planté's battery. Faure's method is still in use today, with only incremental improvements to paste composition, curing, and structure and composition of the grid to which the paste is applied.
The grid developed by Faure was of pure lead with connecting rods of lead at right angles. In contrast, present-day grids are structured for improved mechanical strength and improved current flow. In addition to different grid patterns, modern-day processes also apply one or two thin fiberglass mats over the grid to distribute the weight more evenly. And while Faure had used pure lead for his grids, within a year these had been superseded by lead–antimony alloys to give the structures additional rigidity. However, high-antimony grids have higher hydrogen evolution, and thus greater outgassing and higher maintenance costs. These issues were identified by U. B. Thomas and W. E. Haring at Bell Labs in the 1930s and eventually led to the development of lead–calcium grid alloys in 1935 for standby power batteries on the U.S. telephone network. Related research led to the development of lead–selenium grid alloys in Europe a few years later. Both lead–calcium and lead–selenium grid alloys still add antimony, albeit in much smaller quantities than the older high-antimony grids: lead–calcium grids have 4–6% antimony while lead–selenium grids have 1–2%. These metallurgical improvements give the grid more strength, which allows it to carry more weight, and therefore more active material, and so the plates can be thicker, which in turn contributes to battery lifespan since there is more material available to shed before the battery becomes unusable. High-antimony alloy grids are still used in batteries intended for frequent cycling, e.g. in motor-starting applications where frequent expansion/contraction of the plates need to be compensated for, but where outgassing is not significant since charge currents remain low. Since the 1950s, batteries designed for infrequent cycling applications increasingly have lead–calcium or lead–selenium alloy grids since these have less hydrogen evolution and thus lower maintenance overhead. Lead–calcium alloy grids are cheaper to manufacture, and have a lower self-discharge rate, and lower watering requirements, but have slightly poorer conductivity, are mechanically weaker, and are more strongly subject to corrosion than cells with lead–selenium alloy grids.
The open-circuit effect is a dramatic loss of battery cycle life, which was observed when calcium was substituted for antimony. It is also known as the antimony free effect.