Electric battery

For other uses, see Battery (disambiguation).

Various cells and batteries (top-left to bottom-right): two AA, one D, one handheld ham radio battery, two 9-volt PP3, two AAA, one C, one camcorder battery, one cordless phone battery.

An electrical battery is one or more electrochemical cells that convert stored chemical energy into electrical energy. Since the invention of the first battery (or "voltaic pile") in 1800 by Alessandro Volta, batteries have become a common power source for many household and industrial applications. According to a 2005 estimate, the worldwide battery industry generates US$48 billion in sales each year, with 6% annual growth.

There are two types of batteries: primary batteries (disposable batteries), which are designed to be used once and discarded when they are exhausted, and secondary batteries (rechargeable batteries), which are designed to be recharged and used multiple times. Miniature cells are used to power devices such as hearing aids and wristwatches; larger batteries provide standby power for telephone exchanges or computer data centers.

History

Main article: History of the battery

The symbol for a battery in a circuit diagram. It originated as a schematic drawing of the earliest type of battery, a voltaic pile.

Strictly, a battery is a collection of multiple electrochemical cells, but in popular usage battery often refers to a single cell. The first electrochemical cell was developed by the Italian physicist Alessandro Volta in 1792, and in 1800 he invented the first battery—for him, a "pile" of cells.

The usage of "battery" to describe electrical devices dates to Benjamin Franklin, who in 1748 described multiple Leyden jars (early electrical capacitors) by analogy to a battery of cannons. Thus Franklin's usage to describe multiple Leyden jars predated Volta's use of multiple galvanic cells.. It is speculated, but not established, that several ancient artifacts consisting of copper sheets and iron bars, and known as Baghdad batteries may have been galvanic cells.

Volta's work was stimulated by the Italian anatomist and physiologist Luigi Galvani, who in 1780 noticed that dissected frog's legs would twitch when struck by a spark from a Leyden jar, an external source of electricity. In 1786 he noticed that twitching would occur during lightning storms. After many years Galvani learned how to produce twitching without using any external source of electricity. In 1791 he published a report on "animal electricity." He created an electric circuit consisting of the frog's leg (FL) and two different metals A and B, each metal touching the frog's leg and each other, thus producing the circuit A-FL-B-A-FL-B...etc. In modern terms, the frog's leg served as both the electrolyte and the sensor, and the metals served as electrodes. He noticed that even though the frog was dead, its legs would twitch when he touched them with the metals.

Within a year, Volta realized the frog's moist tissues could be replaced by cardboard soaked in salt water, and the frog's muscular response could be replaced by another form of electrical detection. He already had studied the electrostatic phenomenon of capacitance, which required measurements of electric charge and of electrical potential ("tension"). Building on this experience, Volta was able to detect electric current through his system, also called a Galvanic cell. The terminal voltage of a cell that is not discharging is called its electromotive force (emf), and has the same unit as electrical potential, named (voltage) and measured in volts, in honor of Volta. In 1800, Volta invented the battery by placing many voltaic cells in series, literally piling them one above the other. This voltaic pile gave a greatly enhanced net emf for the combination, with a voltage of about 50 volts for a 32-cell pile. In many parts of Europe batteries continue to be called piles.

Volta did not appreciate that the voltage was due to chemical reactions. He thought that his cells were an inexhaustible source of energy, and that the associated chemical effects (e.g. corrosion) were a mere nuisance, rather than an unavoidable consequence of their operation, as Michael Faraday showed in 1834. According to Faraday, cations (positively charged ions) are attracted to the cathode, and anions (negatively charged ions) are attracted to the anode.

Although early batteries were of great value for experimental purposes, in practice their voltages fluctuated and they could not provide a large current for a sustained period. Later, starting with the Daniell cell in 1836, batteries provided more reliable currents and were adopted by industry for use in stationary devices, particularly in telegraph networks where they were the only practical source of electricity, since electrical distribution networks did not exist at the time. These wet cells used liquid electrolytes, which were prone to leakage and spillage if not handled correctly. Many used glass jars to hold their components, which made them fragile. These characteristics made wet cells unsuitable for portable appliances. Near the end of the nineteenth century, the invention of dry cell batteries, which replaced the liquid electrolyte with a paste, made portable electrical devices practical.

Since then, batteries have gained popularity as they became portable and useful for a variety of purposes.

Principle of operation

Main article: Electrochemical cell

A voltaic cell for demonstration purposes. In this example the two half-cells are linked by a salt bridge separator that permits the transfer of ions, but not water molecules.

A battery is a device that converts chemical energy directly to electrical energy. It consists of a number of voltaic cells; each voltaic cell consists of two half cells connected in series by a conductive electrolyte containing anions and cations. One half-cell includes electrolyte and the electrode to which anions (negatively charged ions) migrate, i.e., the anode or negative electrode; the other half-cell includes electrolyte and the electrode to which cations (positively charged ions) migrate, i.e., the cathode or positive electrode. In the redox reaction that powers the battery, reduction (addition of electrons) occurs to cations at the cathode, while oxidation (removal of electrons) occurs to anions at the anode. The electrodes do not touch each other but are electrically connected by the electrolyte. Many cells use two half-cells with different electrolytes. In that case each half-cell is enclosed in a container, and a separator that is porous to ions, but not the bulk of the electrolytes, prevents mixing.

Each half cell has an electromotive force (or emf), determined by its ability to drive electric current from the interior to the exterior of the cell. The net emf of the cell is the difference between the emfs of its half-cells, as first recognized by Volta. Therefore, if the electrodes have emfs ${\mathcal {E}}_{1}$ and ${\mathcal {E}}_{2}$, then the net emf is ${\mathcal {E}}_{2}-{\mathcal {E}}_{1}$; in other words, the net emf is the difference between the reduction potentials of the half-reactions.

The electrical driving force or $\displaystyle {\Delta V_{bat}}$ across the terminals of a cell is known as the terminal voltage (difference) and is measured in volts. The terminal voltage of a cell that is neither charging nor discharging is called the open-circuit voltage and equals the emf of the cell. Because of internal resistance, the terminal voltage of a cell that is discharging is smaller in magnitude than the open-circuit voltage and the terminal voltage of a cell that is charging exceeds the open-circuit voltage. An ideal cell has negligible internal resistance, so it would maintain a constant terminal voltage of ${\mathcal {E}}$ until exhausted, then dropping to zero. If such a cell maintained 1.5 volts and stored a charge of one coulomb then on complete discharge it would perform 1.5 joule of work. In actual cells, the internal resistance increases under discharge, and the open circuit voltage also decreases under discharge. If the voltage and resistance are plotted against time, the resulting graphs typically are a curve; the shape of the curve varies according to the chemistry and internal arrangement employed.

As stated above, the voltage developed across a cell's terminals depends on the energy release of the chemical reactions of its electrodes and electrolyte. Alkaline and carbon-zinc cells have different chemistries but approximately the same emf of 1.5 volts; likewise NiCd and NiMH cells have different chemistries, but approximately the same emf of 1.2 volts. On the other hand the high electrochemical potential changes in the reactions of lithium compounds give lithium cells emfs of 3 volts or more.

Categories and types of batteries

Main article: List of battery types

From top to bottom: SR41/AG3, SR44/AG13 (button cells), a 9-volt PP3 battery, an AAA cell, an AA cell, a C cell, a D Cell, and a large 3R12. The ruler's unit is in centimeters.

Batteries are classified into two broad categories, each type with advantages and disadvantages.

• Primary batteries irreversibly (within limits of practicality) transform chemical energy to electrical energy. When the initial supply of reactants is exhausted, energy cannot be readily restored to the battery by electrical means.
• Secondary batteries can be recharged; that is, they can have their chemical reactions reversed by supplying electrical energy to the cell, restoring their original composition.

Historically, some types of primary batteries used, for example, for telegraph circuits, were restored to operation by replacing the components of the battery consumed by the chemical reaction. Secondary batteries are not indefinitely rechargeable due to dissipation of the active materials, loss of electrolyte and internal corrosion.

Primary batteries

Main article: Primary cell

Primary batteries can produce current immediately on assembly. Disposable batteries are intended to be used once and discarded. These are most commonly used in portable devices that have low current drain, are only used intermittently, or are used well away from an alternative power source, such as in alarm and communication circuits where other electric power is only intermittently available. Disposable primary cells cannot be reliably recharged, since the chemical reactions are not easily reversible and active materials may not return to their original forms. Battery manufacturers recommend against attempting to recharge primary cells.

Common types of disposable batteries include zinc-carbon batteries and alkaline batteries. Generally, these have higher energy densities than rechargeable batteries, but disposable batteries do not fare well under high-drain applications with loads under 75 ohms (75 Ω).

Secondary batteries

Main article: Rechargeable battery

Secondary batteries must be charged before use; they are usually assembled with active materials in the discharged state. Rechargeable batteries or secondary cells can be recharged by applying electric current, which reverses the chemical reactions that occur during its use. Devices to supply the appropriate current are called chargers or rechargers.

The oldest form of rechargeable battery is the lead-acid battery. This battery is notable in that it contains a liquid in an unsealed container, requiring that the battery be kept upright and the area be well ventilated to ensure safe dispersal of the hydrogen gas produced by these batteries during overcharging. The lead-acid battery is also very heavy for the amount of electrical energy it can supply. Despite this, its low manufacturing cost and its high surge current levels make its use common where a large capacity (over approximately 10Ah) is required or where the weight and ease of handling are not concerns.

A common form of the lead-acid battery is the modern car battery, which can generally deliver a peak current of 450 amperes. An improved type of liquid electrolyte battery is the sealed valve regulated lead acid (VRLA) battery, popular in the automotive industry as a replacement for the lead-acid wet cell. The VRLA battery uses an immobilized sulfuric acid electrolyte, reducing the chance of leakage and extending shelf life. VRLA batteries have the electrolyte immobilized, usually by one of two means:

• Gel batteries (or "gel cell") contain a semi-solid electrolyte to prevent spillage.
• Absorbed Glass Mat (AGM) batteries absorb the electrolyte in a special fiberglass matting.

Other portable rechargeable batteries include several "dry cell" types, which are sealed units and are therefore useful in appliances such as mobile phones and laptop computers. Cells of this type (in order of increasing power density and cost) include nickel-cadmium (NiCd), nickel-zinc (NiZn), nickel metal hydride (NiMH) and lithium-ion (Li-ion) cells. By far, Li-ion has the highest share of the dry cell rechargeable market. Meanwhile, NiMH has replaced NiCd in most applications due to its higher capacity, but NiCd remains in use in power tools, two-way radios, and medical equipment. NiZn is a new technology that is not yet well established commercially.

Recent developments include batteries with embedded functionality such as USBCELL, with a built-in charger and USB connector within the AA format, enabling the battery to be charged by plugging into a USB port without a charger, and low self-discharge (LSD) mix chemistries such as Hybrio, ReCyko, and Eneloop, where cells are precharged prior to shipping.

Battery cell types

There are many general types of electrochemical cells, according to chemical processes applied and design chosen. The variation includes galvanic cells, electrolytic cells, fuel cells, flow cells and voltaic piles.

Wet cell

A wet cell battery has a liquid electrolyte. Other names are flooded cell since the liquid covers all internal parts, or vented cell since gases produced during operation can escape to the air. Wet cells were a precursor to dry cells and are commonly used as a learning tool for electrochemistry. It is often built with common laboratory supplies, like beakers, for demonstrations of how electrochemical cells work. A particular type of wet cell known as a concentration cell is important in understanding corrosion. Wet cells may be primary cells (non-rechargeable) or secondary cells (rechargeable). Originally all practical primary batteries such as the Daniell cell were built as open-topped glass jar wet cells. Other primary wet cells are the Leclanche cell, Grove cell, Bunsen cell, Chromic acid cell, Clark cell and Weston cell. The Leclanche cell chemistry was adapted to the first dry cells.

Wet cells are still used in automobile batteries and in industry for standby power for switchgear, telecommunication or large uninterruptible power supplies, but in many places batteries with gel cells have been used instead. These applications commonly use lead-acid or nickel-cadmium cells.

Dry cell

"Dry cell" redirects here. For the heavy metal band, see Dry Cell (band).

Line art drawing of a dry cell:
1. brass cap, 2. plastic seal, 3. expansion space, 4. porous cardboard, 5. zinc can, 6. carbon rod, 7. chemical mixture.

A dry cell has the electrolyte immobilized as a paste, with only enough moisture in the paste to allow current to flow. As opposed to a wet cell, the battery can be operated in any random position, and will not spill its electrolyte if inverted.

While a dry cell's electrolyte is not truly completely free of moisture and must contain some moisture to function, it has the advantage of containing no sloshing liquid that might leak or drip out when inverted or handled roughly, making it highly suitable for small portable electric devices. By comparison, the first wet cells were typically fragile glass containers with lead rods hanging from the open top, and needed careful handling to avoid spillage. An inverted wet cell would leak, while a dry cell would not. Lead-acid batteries would not achieve the safety and portability of the dry cell until the development of the gel battery.

A common dry cell battery is the zinc-carbon battery, using a cell sometimes called the dry Leclanché cell, with a nominal voltage of 1.5 volts, the same nominal voltage as the alkaline battery (since both use the same zinc-manganese dioxide combination).

The makeup of a standard dry cell is a zinc anode (negative pole), usually in the form of a cylindrical pot, with a carbon cathode (positive pole) in the form of a central rod. The electrolyte is ammonium chloride in the form of a paste next to the zinc anode. The remaining space between the electrolyte and carbon cathode is taken up by a second paste consisting of ammonium chloride and manganese dioxide, the latter acting as a depolariser. In some more modern types of so called 'high power' batteries, the ammonium chloride has been replaced by zinc chloride.

Molten salt

A molten salt battery is a primary or secondary battery that uses a molten salt as its electrolyte. Their energy density and power density makes them potentially useful for electric vehicles, but they must be carefully insulated to retain heat.

Reserve

A reserve battery can be stored for a long period of time and is activated when its internal parts (usually electrolyte) are assembled. For example, a battery for an electronic fuze might be activated by the impact of firing a gun, breaking a capsule of electrolyte to activate the battery and power the fuze's circuits. Reserve batteries are usually designed for a short service life (seconds or minutes) after long storage (years). A water-activated battery for oceanographic instruments or military applications becomes activated on immersion in water.

Battery cell performance

A battery's characteristics may vary over load cycle, charge cycle and over life time due to many factors including internal chemistry, current drain and temperature.

Battery capacity and discharging

A device to check battery voltage.

The more electrolyte and electrode material there is in the cell, the greater the capacity of the cell. Thus a small cell has less capacity than a larger cell, given the same chemistry (e.g. alkaline cells), though they develop the same open-circuit voltage.

Because of the chemical reactions within the cells, the capacity of a battery depends on the discharge conditions such as the magnitude of the current (which may vary with time), the allowable terminal voltage of the battery, temperature and other factors. The available capacity of a battery depends upon the rate at which it is discharged. If a battery is discharged at a relatively high rate, the available capacity will be lower than expected.

The battery capacity that battery manufacturers print on a battery is usually the product of 20 hours multiplied by the maximum constant current that a new battery can supply for 20 hours at 68 F° (20 C°), down to a predetermined terminal voltage per cell. A battery rated at 100 A·h will deliver 5 A over a 20 hour period at room temperature. However, if it is instead discharged at 50 A, it will have a lower apparent capacity.

The relationship between current, discharge time, and capacity for a lead acid battery is approximated (over a certain range of current values) by Peukert's law:

$t={\frac {Q_{P}}{I^{k}}}$

where

$Q_{P}$ is the capacity when discharged at a rate of 1 amp.

$I$ is the current drawn from battery (A).

$t$ is the amount of time (in hours) that a battery can sustain.

$k$ is a constant around 1.3.

For low values of I internal self-discharge must be included.

In practical batteries, internal energy losses, and limited rate of diffusion of ions through the electrolyte, cause the efficiency of a battery to vary at different discharge rates. When discharging at low rate, the battery's energy is delivered more efficiently than at higher discharge rates, but if the rate is too low, it will self-discharge during the long time of operation, again lowering its efficiency.

Installing batteries with different A·h ratings will not affect the operation of a device rated for a specific voltage unless the load limits of the battery are exceeded. High-drain loads like digital cameras can result in lower actual energy, most notably for alkaline batteries. For example, a battery rated at 2000 mA·h would not sustain a current of 1 A for the full two hours, if it had been rated at a 10-hour or 20-hour discharge.

Fastest charging, largest, and lightest batteries

Lithium iron phosphate (LiFePO₄) batteries are the fastest charging and discharging, next to supercapacitors. The world's largest battery is in Fairbanks, Alaska, composed of Ni-Cd cells. Sodium-sulfur batteries are being used to store wind power. Lithium-sulfur batteries have been used on the longest and highest solar powered flight. The speed of recharging for lithium-ion batteries may be increased by manipulation.

Battery lifetime

Life of primary batteries

Even if never taken out of the original package, disposable (or "primary") batteries can lose 8 to 20 percent of their original charge every year at a temperature of about 20°–30°C. This is known as the "self discharge" rate and is due to non-current-producing "side" chemical reactions, which occur within the cell even if no load is applied to it. The rate of the side reactions is reduced if the batteries are stored at low temperature, although some batteries can be damaged by freezing. High or low temperatures may reduce battery performance. This will affect the initial voltage of the battery. For an AA alkaline battery this initial voltage is approximately normally distributed around 1.6 volts.

Discharging performance of all batteries drops at low temperature.

Battery sizes

Main article: List of battery sizes

Lifespan of rechargeable batteries

Rechargeable batteries.

Rechargeable batteries self-discharge more rapidly than disposable alkaline batteries, especially nickel-based batteries; a freshly charged NiCd loses 10% of its charge in the first 24 hours, and thereafter discharges at a rate of about 10% a month. However, modern lithium designs have reduced the self-discharge rate to a relatively low level (but still poorer than for primary batteries). Most nickel-based batteries are partially discharged when purchased, and must be charged before first use.

Although rechargeable batteries have their energy content restored by charging, some deterioration occurs on each charge/discharge cycle. Low-capacity nickel metal hydride (NiMH) batteries (1700-2000 mA·h) can be charged for about 1000 cycles, whereas high capacity NiMH batteries (above 2500 mA·h) can be charged for about 500 cycles. Nickel cadmium (NiCd) batteries tend to be rated for 1,000 cycles before their internal resistance permanently increases beyond usable values. Normally a fast charge, rather than a slow overnight charge, will shorten battery lifespan. However, if the overnight charger is not "smart" and cannot detect when the battery is fully charged, then overcharging is likely, which also damages the battery. Degradation usually occurs because electrolyte migrates away from the electrodes or because active material falls off the electrodes. NiCd batteries suffer the drawback that they should be fully discharged before recharge. Without full discharge, crystals may build up on the electrodes, thus decreasing the active surface area and increasing internal resistance. This decreases battery capacity and causes the "memory effect". These electrode crystals can also penetrate the electrolyte separator, thereby causing shorts. NiMH, although similar in chemistry, does not suffer from memory effect to quite this extent. When a battery reaches the end of its lifetime, it will not suddenly lose all of its capacity; rather, its capacity will gradually decrease.

Automotive lead-acid rechargeable batteries have a much harder life. Because of vibration, shock, heat, cold, and sulfation of their lead plates, few automotive batteries last beyond six years of regular use. Automotive starting batteries have many thin plates to provide as much current as possible in a reasonably small package. In general, the thicker the plates, the longer the life of the battery. Typically they are only drained a small amount before recharge. Care should be taken to avoid deep discharging a starting battery, since each charge and discharge cycle causes active material to be shed from the plates.

"Deep-cycle" lead-acid batteries such as those used in electric golf carts have much thicker plates to aid their longevity. The main benefit of the lead-acid battery is its low cost; the main drawbacks are its large size and weight for a given capacity and voltage. Lead-acid batteries should never be discharged to below 20% of their full capacity, because internal resistance will cause heat and damage when they are recharged. Deep-cycle lead-acid systems often use a low-charge warning light or a low-charge power cut-off switch to prevent the type of damage that will shorten the battery's life.

Extending battery life

Battery life can be extended by storing the batteries at a low temperature, as in a refrigerator or freezer, which slows the chemical reactions in the battery. Such storage can extend the life of alkaline batteries by about 5%, while the charge of rechargeable batteries can be extended from a few days up to several months. To reach their maximum voltage, batteries must be returned to room temperature; discharging an alkaline battery at 250 mAh at 0°C is only half as efficient as it is at 20°C. As a result, alkaline battery manufacturers like Duracell do not recommend refrigerating or freezing batteries.

Prolonging life in multiple cells through cell balancing

Analog front ends that balance cells and eliminate mismatches of cells in series or parallel combination significantly improve battery efficiency and increase the overall pack capacity. As the number of cells and load currents increase, the potential for mismatch also increases. There are two kinds of mismatch in the pack: state-of-charge (SOC) and capacity/energy (C/E) mismatch. Though the SOC mismatch is more common, each problem limits the pack capacity (mAh) to the capacity of the weakest cell.

Cell balancing principle

Battery pack cells are balanced when all the cells in the battery pack meet two conditions:

1. If all cells have the same capacity, then they are balanced when they have the same State of Charge (SOC.) In this case, the Open Circuit Voltage (OCV) is a good measure of the SOC. If, in an out of balance pack, all cells can be differentially charged to full capacity (balanced), then they will subsequently cycle normally without any additional adjustments. This is mostly a one shot fix.
2. If the cells have different capacities, they are also considered balanced when the SOC is the same. But, since SOC is a relative measure, the absolute amount of capacity for each cell is different. To keep the cells with different capacities at the same SOC, cell balancing must provide differential amounts of current to cells in the series string during both charge and discharge on every cycle.

Cell balancing electronics

Cell balancing is defined as the application of differential currents to individual cells (or combinations of cells) in a series string. Normally, of course, cells in a series string receive identical currents. A battery pack requires additional components and circuitry to achieve cell balancing. However, the use of a fully integrated analog front end for cell balancing reduces the required external components to just balancing resistors.

It is important to recognize that the cell mismatch results more from limitations in process control and inspection than from variations inherent in the Lithium Ion chemistry. The use of a fully integrated analog front end for cell balancing can improve the performance of series connected Li-ion Cells by addressing both SOC and C/E issues. SOC mismatch can be remedied by balancing the cell during an initial conditioning period and subsequently only during the charge phase. C/E mismatch remedies are more difficult to implement and harder to measure and require balancing during both charge and discharge periods.

This type of solution eliminates the quantity of external components, as for discrete capacitors, diodes and most other resistors to achieve balance.

Hazards

Explosion

A battery explosion is caused by the misuse or malfunction of a battery, such as attempting to recharge a primary (non-rechargeable) battery, or short circuiting a battery. With car batteries, explosions are most likely to occur when a short circuit generates very large currents. In addition, car batteries liberate hydrogen when they are overcharged (because of electrolysis of the water in the electrolyte). Normally the amount of overcharging is very small, as is the amount of explosive gas developed, and the gas dissipates quickly. However, when "jumping" a car battery, the high current can cause the rapid release of large volumes of hydrogen, which can be ignited by a nearby spark (for example, when removing the jumper cables).

When a battery is recharged at an excessive rate, an explosive gas mixture of hydrogen and oxygen may be produced faster than it can escape from within the walls of the battery, leading to pressure build-up and the possibility of the battery case bursting. In extreme cases, the battery acid may spray violently from the casing of the battery and cause injury. Overcharging—that is, attempting to charge a battery beyond its electrical capacity—can also lead to a battery explosion, leakage, or irreversible damage to the battery. It may also cause damage to the charger or device in which the overcharged battery is later used. Additionally, disposing of a battery in fire may cause an explosion as steam builds up within the sealed case of the battery.

Leakage

Leaked alkaline battery.

Many battery chemicals are corrosive, poisonous, or both. If leakage occurs, either spontaneously or through accident, the chemicals released may be dangerous.

For example, disposable batteries often use a zinc "can" as both a reactant and as the container to hold the other reagents. If this kind of battery is run all the way down, or if it is recharged after running down too far, the reagents can emerge through the cardboard and plastic that form the remainder of the container. The active chemical leakage can then damage the equipment that the batteries were inserted into. For this reason, many electronic device manufacturers recommend removing the batteries from devices that will not be used for extended periods of time.

Environmental concerns

The widespread use of batteries has created many environmental concerns, such as toxic metal pollution. Battery manufacture consumes resources and often involves hazardous chemicals. Used batteries also contribute to electronic waste. Some areas now have battery recycling services available to recover some of the materials from used batteries. Batteries may be harmful or fatal if swallowed. Recycling or proper disposal prevents dangerous elements (such as lead, mercury, and cadmium) found in some types of batteries from entering the environment. In the United States, Americans purchase nearly three billion batteries annually, and about 179,000 tons of those end up in landfills across the country.

In the United States, the Mercury-Containing and Rechargeable Battery Management Act of 1996 banned the sale of mercury-containing batteries (except small button cell batteries), enacted uniform labeling requirements for rechargeable batteries, and required that rechargeable batteries be easily removable. California, and New York City prohibit the disposal of rechargeable batteries in solid waste, and along with Maine require recycling of cell phones. The rechargeable battery industry has nationwide recycling programs in the United States and Canada, with dropoff points at local retailers.

The Battery Directive of the European Union has similar requirements, in addition to requiring increased recycling of batteries, and promoting research on improved battery recycling methods.

Ingestion

Small button/disk batteries can be swallowed by young children. While in the digestive tract the battery's electrical discharge can burn the tissues and can be serious enough to lead to death. Disk batteries do not usually cause problems unless they become lodged in the gastrointestinal (GI) tract. The most common place disk batteries become lodged, resulting in clinical sequelae, is the esophagus. Batteries that successfully traverse the esophagus are unlikely to lodge at any other location. The likelihood that a disk battery will lodge in the esophagus is a function of the patient's age and the size of the battery. Disk batteries of 16 mm have become lodged in the esophagi of 2 children younger than 1 year. Older children do not have problems with batteries smaller than 21-23 mm. For comparison, a dime is 18 mm, a nickel is 21 mm, and a quarter is 24 mm. Liquefaction necrosis may occur because sodium hydroxide is generated by the current produced by the battery (usually at the anode). Perforation has occurred as rapidly as 6 hours after ingestion.

Battery chemistry

Primary battery chemistries

(includes data from the energy density article)

Chemistry	Nominal Cell Voltage	Specific Energy [MJ/kg]	Elaboration
Zinc–carbon	1.5	0.13	Inexpensive.
Zinc–chloride	1.5		Also known as "heavy duty", inexpensive.
Alkaline (zinc–manganese dioxide)	1.5	0.4-0.59	Moderate energy density. Good for high and low drain uses.
Oxy nickel hydroxide (zinc–manganese dioxide/oxy nickel hydroxide)	1.7		Moderate energy density. Good for high drain uses
Lithium (lithium–copper oxide) Li–CuO	1.7		No longer manufactured. Replaced by silver oxide (IEC-type "SR") batteries.
Lithium (lithium–iron disulfide) LiFeS2	1.5		Expensive. Used in 'plus' or 'extra' batteries.
Lithium (lithium–manganese dioxide) LiMnO2	3.0	0.83-1.01	Expensive. Only used in high-drain devices or for long shelf life due to very low rate of self discharge. 'Lithium' alone usually refers to this type of chemistry.
Mercury oxide	1.35		High drain and constant voltage. Banned in most countries because of health concerns.
Zinc–air	1.35–1.65	1.59	Mostly used in hearing aids.
Silver-oxide (silver–zinc)	1.55	0.47	Very expensive. Only used commercially in 'button' cells.

Rechargeable battery chemistries

(includes data from energy density article)

Chemistry	Cell Voltage	Specific Energy [MJ/kg]	Comments
NiCd	1.2	0.14	Inexpensive. High/low drain, moderate energy density. Can withstand very high discharge rates with virtually no loss of capacity. Moderate rate of self discharge. Reputed to suffer from memory effect (which is alleged to cause early failure). Environmental hazard due to Cadmium - use now virtually prohibited in Europe.
Lead acid	2.1	0.14	Moderately expensive. Moderate energy density. Moderate rate of self discharge. Higher discharge rates result in considerable loss of capacity. Does not suffer from memory effect. Environmental hazard due to Lead. Common use - Automobile batteries
NiMH	1.2	0.36	Inexpensive. Performs better than alkaline batteries in higher drain devices. Traditional chemistry has high energy density, but also a high rate of self-discharge. Newer chemistry has low self-discharge rate, but also a ~25% lower energy density. Very heavy. Used in some cars.
NiZn	1.6	0.36	Moderately inexpensive. High drain device suitable. Low self-discharge rate. Voltage closer to alkaline primary cells than other secondary cells. No toxic components. Newly introduced to the market (2009). Has not yet established a track record. Limited size availability.
Lithium ion	3.6	0.46	Very expensive. Very high energy density. Not usually available in "common" battery sizes (but see RCR-V3 for a counter-example). Very common in laptop computers, moderate to high-end digital cameras and camcorders, and cellphones. Very low rate of self discharge. Volatile: Chance of explosion if short circuited, allowed to overheat, or not manufactured with rigorous quality standards.

Homemade cells

Almost any liquid or moist object that has enough ions to be electrically conductive can serve as the electrolyte for a cell. As a novelty or science demonstration, it is possible to insert two electrodes made of different metals into a lemon, potato, etc. and generate small amounts of electricity. "Two-potato clocks" are also widely available in hobby and toy stores; they consist of a pair of cells, each consisting of a potato (lemon, et cetera) with two electrodes inserted into it, wired in series to form a battery with enough voltage to power a digital clock. Homemade cells of this kind are of no real practical use, because they produce far less current—and cost far more per unit of energy generated—than commercial cells, due to the need for frequent replacement of the fruit or vegetable. In addition, one can make a voltaic pile from two coins (such as a nickel and a penny) and a piece of paper towel dipped in salt water. Such a pile would make very little voltage itself, but when many of them are stacked together in series, they can replace normal batteries for a short amount of time.

Sony has developed a biological battery that generates electricity from sugar in a way that is similar to the processes observed in living organisms. The battery generates electricity through the use of enzymes that break down carbohydrates, which are essentially sugar. A similarly designed sugar drink powered phone using enzymes to generate electricity from carbohydrates that covers the phone’s electrical needs. It only needs a pack of sugary drink and it generates water and oxygen while the battery dies out.

Lead acid cells can easily be manufactured at home, but a tedious charge/discharge cycle is needed to 'form' the plates. This is a process whereby lead sulfate forms on the plates, and during charge is converted to lead dioxide (positive plate) and pure lead (negative plate). Repeating this process results in a microscopically rough surface, with far greater surface area being exposed. This increases the current the cell can deliver. For an example, see [3].

Daniell cells are also easy to make at home. Aluminum-air batteries can also be produced with high purity aluminum. Aluminum foil batteries will produce some electricity, but they are not very efficient, in part because a significant amount of hydrogen gas is produced.

See also

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• Battery electric vehicle
• Battery (vacuum tube)
• Battery Directive
• Battery holder
• Battery isolator
• Battery Management System (BMS)
• Battery nomenclature
• Battery pack
• Battery recycling
• Battery terminals
• Depth Of Discharge (DOD)
• Electrochemical cell

• Energy density
• Energy storage
• Flexible battery
• Galvanic cell
• List of battery sizes
• List of battery types
• Magnetic battery
• Nano titanate
• Nanowire battery
• Printed battery
• Rechargeable battery
• State Of Charge (SOC)
• State Of Health (SOH)
• Thermal runaway
• Trickle charging

Further reading

• Dingrando, Laurel (2007). Chemistry: Matter and Change. New York: Glencoe/McGraw-Hill. ISBN 978-0-07-877237-5. {{cite book}}: Unknown parameter |coauthors= ignored (|author= suggested) (help) Ch. 21 (pp. 662–695) is on electrochemistry.
• Fink, Donald G. (1978). Standard Handbook for Electrical Engineers, Eleventh Edition. New York: McGraw-Hill. ISBN 0-07020974-X. {{cite book}}: Unknown parameter |coauthors= ignored (|author= suggested) (help)
• Knight, Randall D. (2004). Physics for Scientists and Engineers: A Strategic Approach. San Francisco: Pearson Education. ISBN 0-8053-8960-1. Chs. 28-31 (pp. 879–995) contain information on electric potential.
• Linden, David (2001). Handbook Of Batteries. New York: McGraw-Hill. ISBN 0-0713-5978-8. {{cite book}}: Unknown parameter |coauthors= ignored (|author= suggested) (help)
• Saslow, Wayne M. (2002). Electricity, Magnetism, and Light. Toronto: Thomson Learning. ISBN 0-12-619455-6. Chs. 8-9 (pp. 336–418) have more information on batteries.

External links

• Professional Battery StoreUS
• Battery University
• The Electropaedia on Battery Chemistries and How They Work
• Non-rechargeable batteries
• HowStuffWorks: How batteries work
• Comprehensive knowledge base about battery technology, battery applications, chargers and ancillary equipment
• DoITPoMS Teaching and Learning Package- "Batteries"