# Battery (electricity)

(Redirected from Dry cell)
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.

In electricity, a battery is a device consisting of one or more electrochemical cells that convert stored chemical energy into electrical energy.[1] Since the invention of the first battery (or "voltaic pile") in 1800 by Alessandro Volta and especially since the technically improved Daniell cell in 1836, 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,[2] with 6% annual growth.[3]

There are two types of batteries: primary batteries (disposable batteries), which are designed to be used once and discarded, and secondary batteries (rechargeable batteries), which are designed to be recharged and used multiple times. Batteries come in many sizes, from miniature cells used to power hearing aids and wristwatches to battery banks the size of rooms that provide standby power for telephone exchanges and computer data centers.

## History

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

In strict terms, a battery is a collection of multiple electrochemical cells, but in popular usage battery often refers to a single cell.[1] For example, a 1.5-volt AAA battery is a single 1.5-volt cell, and a 9-volt battery has six 1.5-volt cells in series. The first electrochemical cell was developed by the Italian physicist Alessandro Volta in 1792, and in 1800 he invented the first battery, a "pile" of many cells in series.[4]

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.[5] Thus Franklin's usage to describe multiple Leyden jars predated Volta's use of multiple galvanic cells.[6] 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.[7]

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.[8] In 1786 he noticed that twitching would occur during lightning storms.[9] 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."[10] 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, piling them one above the other. This voltaic pile gave a greatly enhanced net emf for the combination,[11] with a voltage of about 50 volts for a 32-cell pile.[12] In many parts of Europe batteries continue to be called piles.[13][14]

Volta did not appreciate that the voltage was due to chemical reactions. He thought that his cells were an inexhaustible source of energy,[15] and that the associated corrosion effects at the electrodes were a mere nuisance, rather than an unavoidable consequence of their operation, as Michael Faraday showed in 1834.[16] According to Faraday, cations (positively charged ions) are attracted to the cathode,[17] and anions (negatively charged ions) are attracted to the anode.[18]

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, in particular in telegraph networks where they were the only practical source of electricity, since electrical distribution networks did not exist at the time.[19] 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.[20]

Since then, batteries have gained popularity as they became portable and useful for a variety of purposes.[21]

## Principle of operation

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.[22] 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, cations are reduced (electrons are added) at the cathode, while anions are oxidized (electrons are removed) at the anode.[23] The electrodes do not touch each other but are electrically connected by the electrolyte. Some cells use two half-cells with different electrolytes. A separator between half-cells allows ions to flow, but prevents mixing of the electrolytes.

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.[12] 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.[24]

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.[25] 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,[26] 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.[27] 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.[25] In actual cells, the internal resistance increases under discharge,[26] 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.[28]

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 zinc–carbon 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.[29] On the other hand the high electrochemical potential changes in the reactions of lithium compounds give lithium cells emfs of 3 volts or more.[30]

## Categories and types of batteries

From top to bottom: a large 4.5-volt (3R12) battery, a D Cell, a C cell, an AA cell, an AAA cell, an AAAA cell, an A23 battery, a 9-volt PP3 battery, and a pair of button cells (CR2032 and LR44).

• 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.[32]
• 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.[33]

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.[34] Secondary batteries are not indefinitely rechargeable due to dissipation of the active materials, loss of electrolyte and internal corrosion.

### Primary batteries

Primary batteries, or primary cells, 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 used only 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.[35]

Common types of disposable batteries include zinc–carbon batteries and alkaline batteries. In general, these have higher energy densities than rechargeable batteries,[36] but disposable batteries do not fare well under high-drain applications with loads under 75 ohms (75 Ω).[31]

### Secondary batteries

Secondary batteries, also known as secondary cells, or rechargeable batteries, must be charged before use; they are usually assembled with active materials in the discharged state. Rechargeable batteries are (re)charged by applying electric current, which reverses the chemical reactions that occur during discharge/use. Devices to supply the appropriate current are called chargers.

The oldest form of rechargeable battery is the lead–acid battery.[37] This technology contains liquid electrolyte 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 it produces during overcharging. The lead–acid battery is also relatively heavy for the amount of electrical energy it can supply. Its low manufacturing cost and its high surge current levels make it common where its capacity (over approximately 10 Ah) is more important than weight and handling issues. A common application is the modern car battery, which can, in general, deliver a peak current of 450 amperes.

The sealed valve regulated lead–acid battery (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.[38] VRLA batteries immobilize the electrolyte, usually by one of two means:

Other portable rechargeable batteries include several sealed "dry cell" types, that are useful in applications 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.[39] Li-ion has by far the highest share of the dry cell rechargeable market.[3] 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.[3] NiZn is a new technology that is not yet well established commercially.

Recent developments include batteries with embedded electronics such as USBCELL, which allows charging an AA cell through a USB connector,[40] and smart battery packs with state-of-charge monitors and battery protection circuits that prevent damage on over-discharge. Low self-discharge (LSD) allows secondary cells to be charged 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.[41]

#### 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, such as 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

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 it to allow current to flow. Unlike a wet cell, a dry cell can operate in any orientation without spilling as it contains no free liquid, making it suitable for portable equipment. 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. Lead–acid batteries did 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 as the alkaline battery (since both use the same zincmanganese dioxide combination).

A standard dry cell comprises 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 (with much lower capacity than standard alkaline batteries), the ammonium chloride is replaced by zinc chloride.

#### Molten salt

Molten salt batteries are primary or secondary batteries that use a molten salt as electrolyte. Their energy density and power density give them potential for use in electric vehicles, but they operate at high temperatures and must be well insulated to retain heat.

#### Reserve

A reserve battery is stored in unassembled form and is activated, ready-charged, when its internal parts are assembled, e.g. by adding electrolyte; it can be stored unactivated for a long period of time. 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, over charge cycle, and over lifetime due to many factors including internal chemistry, current drain, and temperature.

## Capacity and discharging

A device to check battery voltage

A battery's capacity is the amount of electric charge it can store. The more electrolyte and electrode material there is in the cell the greater the capacity of the cell. A small cell has less capacity than a larger cell with the same chemistry, and they develop the same open-circuit voltage.[42]

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.[42] The available capacity of a battery depends upon the rate at which it is discharged.[43] If a battery is discharged at a relatively high rate, the available capacity will be lower than expected.

The capacity printed on a battery is usually the product of 20 hours multiplied by the constant current that a new battery can supply for 20 hours at 68 F° (20 C°), down to a specified 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 discharged at 50 A, it will have a lower capacity.[44]

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.

Internal energy losses and limited rate of diffusion of ions through the electrolyte cause the efficiency of a real 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,[44] but if the rate is very low, it will partly 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 (except for the time it will work for) rated for a specific voltage unless the load limits of the battery are exceeded. High-drain loads such as digital cameras can result in delivery of less total energy, as happens with alkaline batteries.[31] For example, a battery rated at 2000 mAh for a 10- or 20-hour discharge would not sustain a current of 1 A for a full two hours as its stated capacity implies.

### C rates

The C-rate signifies a discharge rate relative to the capacity of a battery in one hour.[45] A rate of 1C would mean an entire 1.6Ah battery would be discharged in 1 hour at a discharge current of 1.6A. A 2C rate would mean a discharge current of 3.2A

### Fastest charging, largest, and lightest batteries

As of 2012 Lithium iron phosphate (LiFePO4) batteries were the fastest-charging and discharging batteries (supercapacitors, in some ways comparable to batteries, charge faster).[46] The world's largest battery, composed of Ni–Cd cells, was in Fairbanks, Alaska.[47] Sodium–sulfur batteries were being used to store wind power.[48] Lithium–sulfur batteries have been used on the longest and highest solar-powered flight.[49] The speed of recharging of lithium-ion batteries can be increased by manufacturing changes.[50]

### Primary batteries

Disposable (or "primary") batteries typically lose 8 to 20 percent of their original charge every year at room temperature (20°–30°C).[51] 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. The rate of the side reactions is reduced if the batteries are stored at lower temperature, although some batteries can be damaged by freezing. High or low working 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.

In contrast to most of today's batteries, the Zamboni pile, invented in 1812, can have a very long service life without refurbishment or recharge, although it supplies current only in the nanoamp range. The Oxford Electric Bell has been ringing almost continuously since 1840 on its original pair of batteries, thought to be Zamboni piles.

Discharging performance of all batteries drops at low temperature.[52]

### Secondary batteries

Storage life of secondary batteries is limited by chemical reactions that occur between the battery parts and the electrolyte; these are called "side reactions". Internal parts may corrode and fail, or the active materials may be slowly converted to inactive forms. Since the active material on the battery plates changes chemical composition on each charge and discharge cycle, active material may be lost due to physical changes of volume; this may limit the cycle life of the battery.

Rechargeable batteries.

Old chemistry rechargeable batteries self-discharge more rapidly than disposable alkaline batteries, especially nickel-based batteries; a freshly charged nickel cadmium (NiCd) battery loses 10% of its charge in the first 24 hours, and thereafter discharges at a rate of about 10% a month.[53] However, newer low self-discharge nickel metal hydride (NiMH) batteries and modern lithium designs have reduced the self-discharge rate to a relatively low level (but still poorer than for primary batteries).[53] Most nickel-based batteries are partially discharged when purchased, and must be charged before first use.[54] Newer NiMH batteries are ready to be used when purchased, and have only 15% discharge in a year.[55]

Although rechargeable batteries have their energy content restored by charging, some deterioration occurs on each charge–discharge cycle. Low-capacity 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.[56] NiCd batteries tend to be rated for 1000 cycles before their internal resistance permanently increases beyond usable values. Under normal circumstances, a fast charge, rather than a slow overnight charge, will shorten battery lifespan.[56] Also, 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.[57] 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.[58] A battery does not suddenly stop working; its capacity gradually decreases over its lifetime, until it can no longer hold sufficient charge.[59]

An analog camcorder battery [lithium ion].

Automotive lead–acid rechargeable batteries have a much harder life.[60] Because of vibration, shock, heat, cold, and sulfation of their lead plates, few automotive batteries last beyond six years of regular use.[61] Automotive starting (SLI: Starting, Lighting, Ignition) 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.[60] They are typically drained only 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.[62] 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.[60] Lead–acid batteries should never be discharged to below 20% of their full capacity,[63] 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.[64]

### 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%; rechargeable batteries can hold their charge much longer, depending upon type.[65] To reach their maximum voltage, batteries must be returned to room temperature; discharging an alkaline battery at 250 mA at 0°C is only half as efficient as it is at 20°C.[36] Alkaline battery manufacturers such as Duracell do not recommend refrigerating batteries.[35]

## Battery sizes

Primary batteries readily available to consumers range from tiny button cells used for electric watches, to the No. 6 cell used for signal circuits or other long duration applications. Secondary cells are made in very large sizes; very large batteries can power a submarine. Large secondary batteries have been used to stabilize the electrical grid and help level out peak loads.

## 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. Car batteries are most likely to explode when a short-circuit generates very large currents. Car batteries liberate hydrogen, which is very explosive, when they are overcharged (because of electrolysis of the water in the electrolyte). The amount of overcharging is usually very small and generates little hydrogen, which 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 explosively by a nearby spark, for example, when disconnecting a jumper cable.

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 bursting of the battery case. 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, in addition to leakage or irreversible damage. It may also cause damage to the charger or device in which the overcharged battery is later used. In addition, 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" both as 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.[66] 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.[67] Batteries may be harmful or fatal if swallowed.[68] 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.[69]

In the United States, the Mercury-Containing and Rechargeable Battery Management Act of 1996 banned the sale of mercury-containing batteries, enacted uniform labeling requirements for rechargeable batteries, and required that rechargeable batteries be easily removable.[70] California, and New York City prohibit the disposal of rechargeable batteries in solid waste, and along with Maine require recycling of cell phones.[71] The rechargeable battery industry has nationwide recycling programs in the United States and Canada, with dropoff points at local retailers.[71]

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.[72] In accordance with this directive all batteries to be sold within the EU must be marked with the "collection symbol" (A crossed out wheeled bin). This must cover at least 3% of the surface of prismatic batteries and 1.5% of the surface of cylindrical batteries. All packaging must be marked likewise.[73]

### Ingestion

Small button cells can be swallowed, particularly by young children. While in the digestive tract the battery's electrical discharge may lead to tissue damage;[74] such damage is occasionally serious and very rarely even leads to death. Ingested 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.[citation needed] Older children do not have problems with batteries smaller than 21–23 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.[75]

## 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.
Nickel oxyhydroxide
(zinc–manganese dioxide/nickel oxyhydroxide)
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[76] 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]
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.
Environmental hazard due to Cadmium – use now virtually prohibited in Europe.
Moderate energy density.
Moderate rate of self discharge.
Higher discharge rates result in considerable loss of capacity.
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.
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.
AgZn 1.86
1.5
0.46 Smaller volume than equivalent Li-ion.
Historically extremely expensive.
Very high energy density.
Very high drain capable.
For many years considered obsolete due to high silver prices.
Cell suffers from oxidation if unused.
Reactions are not fully understood.
Terminal voltage very stable but suddenly drops to 1.5 volts at 70-80% charge (believed to be
due to presence of both argentous and argentic oxide in positive plate - one is consumed first).
Has been used in lieu of primary battery (moon buggy).
Is being developed once again as a replacement for Li-ion.
Lithium ion 3.6 0.46 Very expensive.
Very high energy density.
Not usually available in "common" battery sizes.
Very common in laptop computers, moderate to high-end digital cameras, camcorders and cellphones.
Very low rate of self discharge.
Terminal voltage unstable (varies from 4.2 to 3.0 volts during discharge).
Volatile: Chance of explosion if short circuited, allowed to overheat, or not manufactured with rigorous quality standards.

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,[77] potato,[78] 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.[79] 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

A voltaic pile can be made from two coins (such as a nickel and a penny) and a piece of paper towel dipped in salt water. Such a pile generates a very low voltage but, when many are stacked in series, they can replace normal batteries for a short time.[80]

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, in essence, sugar.[81]

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 in which 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. Aluminium–air batteries can also be produced with high-purity aluminium. Aluminium foil batteries will produce some electricity, but are not very efficient, in part because a significant amount of hydrogen gas is produced.

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