An electrical battery, first named by Benjamin Franklin in 1748, is a combination of two or more electrochemical cells used to convert stored chemical energy into electrical energy. Since the invention of the first voltaic pile in 1800 by Alessandro Volta, the battery has 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 (UPS) for telephone exchanges or computer data centers

How Batteries Work

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.

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

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).

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.

Rechargeable battery chemistries

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 and UPS applications.
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.
Lithiumion 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 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.
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