Alcohol

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This article is about the generic chemistry term. For the kind of alcohol found in alcoholic beverages, see Ethanol. For beverages containing alcohol, see Alcoholic beverage. For other uses, see Alcohol (disambiguation).

The hydroxyl (OH) functional group in an alcohol molecule

In chemistry, an alcohol is any organic compound in which a hydroxyl functional group (-OH) is bound to a carbon atom, usually connected to other carbon or hydrogen atoms.
An important class are the simple acyclic alcohols, the general formula for which is C_nH_2n+1OH. Of those, ethanol (C₂H₅OH) is the type of alcohol found in alcoholic beverages, and in common speech the word alcohol refers specifically to ethanol.
Other alcohols are usually described with a clarifying adjective, as in isopropyl alcohol (propan-2-ol) or wood alcohol (methyl alcohol, or methanol). The suffix -ol appears in the IUPAC chemical name of all substances where the hydroxyl group is the functional group with the highest priority; in substances where a higher priority group is present the prefix hydroxy- will appear in the IUPAC name. The suffix -ol in non-systematic names (such as paracetamol or cholesterol) also typically indicates that the substance includes a hydroxyl functional group and so can be termed an alcohol, but many substances (such as citric acid, lactic acid or sucrose) contain one or more hydroxyl functional groups without using the suffix.

Simple alcohols

The most commonly used alcohol is ethanol, C₂H₅OH, with the ethane backbone. Ethanol has been produced and consumed by humans for millennia, in the form of fermented and distilled alcoholic beverages. It is a clear flammable liquid that boils at 78.4 °C, which is used as an industrial solvent, car fuel, and raw material in the chemical industry. In the US and some other countries, because of legal and tax restrictions on alcohol consumption, ethanol destined for other uses often contains additives that make it unpalatable (such as Bitrex) or poisonous (such as methanol). Ethanol in this form is known generally as denatured alcohol; when methanol is used, it may be referred to as methylated spirits ("Meths") or "surgical spirits".
The simplest alcohol is methanol, CH₃OH, which was formerly obtained by the distillation of wood and therefore is called "wood alcohol". It is a clear liquid resembling ethanol in smell and properties, with a slightly lower boiling point (64.7 °C), and is used mainly as a solvent, fuel, and raw material. Unlike ethanol, methanol is extremely toxic: one sip (as little as 10 ml) can cause permanent blindness by destruction of the optic nerve and 30 ml (one fluid ounce) is potentially fatal.^[1]
Two other alcohols whose uses are relatively widespread (though not so much as those of methanol and ethanol) are propanol and butanol. Like ethanol, they can be produced by fermentation processes. (However, the fermenting agent is a bacterium, Clostridium acetobutylicum, that feeds on cellulose, not sugars like the Saccharomyces yeast that produces ethanol.) Saccharomyces yeast are known to produce these higher alcohols at temperatures above 75 F. These alcohols are called fusel alcohols or fusel oils in brewing and tend to have a spicy or peppery flavor. They are considered a fault in most styles of beer.^{[citation needed]}
Simple alcohols, particularly ethanol and methanol, possess denaturing and inert rendering properties, leading to their use as anti-microbial agents in medicine, pharmacy and industry.^{[citation needed]}

Nomenclature

Systematic names

In the IUPAC system, the name of the alkane chain loses the terminal "e" and adds "ol", e.g. "methanol" and "ethanol".^[2] When necessary, the position of the hydroxyl group is indicated by a number between the alkane name and the "ol": propan-1-ol for CH₃CH₂CH₂OH, propan-2-ol for CH₃CH(OH)CH₃. Sometimes, the position number is written before the IUPAC name: 1-propanol and 2-propanol. If a higher priority group is present (such as an aldehyde, ketone or carboxylic acid), then it is necessary to use the prefix "hydroxy",^[2] for example: 1-hydroxy-2-propanone (CH₃COCH₂OH).

Some examples of simple alcohols and how to name them

The IUPAC nomenclature is used in scientific publications and where precise identification of the substance is important. In other less formal contexts, an alcohol is often called with the name of the corresponding alkyl group followed by the word "alcohol", e.g. methyl alcohol, ethyl alcohol. Propyl alcohol may be n-propyl alcohol or isopropyl alcohol depending on whether the hydroxyl group is bonded to the 1st or 2nd carbon on the propane chain.
Alcohols are classified into primary, secondary and tertiary, based upon the number of carbon atoms connected to the carbon atom that bears the hydroxyl group. Namely, the primary alcohols have general formulas RCH₂OH; secondary ones are RR'CHOH; and tertiary ones are RR'R"COH, where R, R' and R" stand for alkyl groups. Ethanol and n-propyl alcohol are primary alcohols; isopropyl alcohol is a secondary one. The prefixes sec- (or s-) and tert- (or t-), conventionally in italics, may be used before the alkyl group's name to distinguish secondary and tertiary alcohols, respectively, from the primary one. For example, isopropyl alcohol is occasionally called sec-propyl alcohol, and the tertiary alcohol (CH₃)₃COH, or 2-methylpropan-2-ol in IUPAC nomenclature, is commonly known as tert-butyl alcohol or tert-butanol.

Common Names

Chemical Formula	IUPAC Name	Common Name
Monohydric alcohols
CH3OH	Methanol	Wood alcohol
C2H5OH	Ethanol	Grain alcohol
C5H11OH	Pentanol	Amyl alcohol
C16H33OH	Hexadecan-1-ol	Cetyl alcohol
Polyhydric alcohols
C2H4(OH)2	Ethane-1 ,2-diol	Ethylene glycol
C3H5(OH)3	Propane-1 ,2,3-triol	Glycerin
C4H6(OH)4	Butane-1 ,2,3,4-tetraol	Erythritol
C5H7(OH)5	Pentane-1 ,2,3,4,5-pentol	Xylitol
C6H8(OH)6	Hexane-1 ,2,3,4,5,6-hexol	Mannitol, Sorbitol
C7H9(OH)7	Heptane-1 ,2,3,4,5,6,7-heptol	Volemitol
Unsaturated aliphatic alcohols
C3H5OH	Prop-2-ene-1-ol	Allyl alcohol
C10H17OH	3,7-Dimethylocta-2,6-dien-1-ol	Geraniol
C3H3OH	Prop-2-in-1-ol	Propargyl alcohol
Alicyclic alcohols
C6H6(OH)6	Cyclohexane-1 ,2,3,4,5,6-geksol	Inositol
C10H19OH	2 - (2-propyl)-5-methyl-cyclohexane-1-ol	Menthol

Etymology

Look up alcohol in Wiktionary, the free dictionary.

The word alcohol appears in English in the 16th century, loaned via French from medical Latin, ultimately from the Arabic الكحل (al-kuḥl, "the kohl, a powder used as an eyeliner").
ال al is Arabic for the definitive article, the in English.
The current Arabic name for alcohol is الكحول al-kuḥūl, re-introduced from western usage.
kuḥl was the name given to the very fine powder, produced by the sublimation of the natural mineral stibnite to form antimony sulfide Sb₂S₃ (hence the essence or "spirit" of the substance), which was used as an antiseptic and eyeliner.
Bartholomew Traheron in his 1543 translation of John of Vigo introduces the word as a term used by "barbarous" (Moorish) authors for "fine powder":

the barbarous auctours use alcohol, or (as I fynde it sometymes wryten) alcofoll, for moost fine poudre.

William Johnson in his 1657 Lexicon Chymicum glosses the word as antimonium sive stibium. By extension, the word came to refer to any fluid obtained by distillation, including "alcohol of wine", the distilled essence of wine. Libavius in Alchymia (1594) has vini alcohol vel vinum alcalisatum. Johnson (1657) glosses alcohol vini as quando omnis superfluitas vini a vino separatur, ita ut accensum ardeat donec totum consumatur, nihilque fæcum aut phlegmatis in fundo remaneat. The word's meaning became restricted to "spirit of wine" (ethanol) in the 18th century, and was again extended to the family of substances so called in modern chemistry from 1850.

Physical and chemical properties

Alcohols have an odor that is often described as “biting” and as “hanging” in the nasal passages.
The hydroxyl group generally makes the alcohol molecule polar. Those groups can form hydrogen bonds to one another and to other compounds (except in certain large molecules where the hydroxyl is protected by steric hindrance of adjacent groups^[3]). This hydrogen bonding means that alcohols can be used as protic solvents. Two opposing solubility trends in alcohols are: the tendency of the polar OH to promote solubility in water, and of the carbon chain to resist it. Thus, methanol, ethanol, and propanol are miscible in water because the hydroxyl group wins out over the short carbon chain. Butanol, with a four-carbon chain, is moderately soluble because of a balance between the two trends. Alcohols of five or more carbons (Pentanol and higher) are effectively insoluble in water because of the hydrocarbon chain's dominance. All simple alcohols are miscible in organic solvents.
Because of hydrogen bonding, alcohols tend to have higher boiling points than comparable hydrocarbons and ethers. The boiling point of the alcohol ethanol is 78.29 °C, compared to 69 °C for the hydrocarbon Hexane (a common constituent of gasoline), and 34.6 °C for Diethyl ether.
Alcohols, like water, can show either acidic or basic properties at the O-H group. With a pK_a of around 16-19 they are generally slightly weaker acids than water, but they are still able to react with strong bases such as sodium hydride or reactive metals such as sodium. The salts that result are called alkoxides, with the general formula RO^- M⁺.
Meanwhile the oxygen atom has lone pairs of nonbonded electrons that render it weakly basic in the presence of strong acids such as sulfuric acid. For example, with methanol:

Alcohols can also undergo oxidation to give aldehydes, ketones or carboxylic acids, or they can be dehydrated to alkenes. They can react to form ester compounds, and they can (if activated first) undergo nucleophilic substitution reactions. The lone pairs of electrons on the oxygen of the hydroxyl group also makes alcohols nucleophiles. For more details see the reactions of alcohols section below.
As one moves from primary to secondary to tertiary alcohols with the same backbone, the hydrogen bond strength, the boiling point,and the acidity typically decrease.

Applications

Total recorded alcohol per capita consumption (15+), in litres of pure alcohol^[4]

Alcohols can be used as a beverage (ethanol only), as fuel and for many scientific, medical, and industrial utilities. Ethanol in the form of alcoholic beverages has been consumed by humans since pre-historic times. A 50% v/v solution of ethylene glycol in water is commonly used as an antifreeze.
Some alcohols, mainly ethanol and methanol, can be used as an alcohol fuel. Fuel performance can be increased in forced induction internal combustion engines by injecting alcohol into the air intake after the turbocharger or supercharger has pressurized the air. This cools the pressurized air, providing a denser air charge, which allows for more fuel, and therefore more power.
Alcohols have applications in industry and science as reagents or solvents. Because of its low toxicity and ability to dissolve non-polar substances, ethanol can be used as a solvent in medical drugs, perfumes, and vegetable essences such as vanilla. In organic synthesis, alcohols serve as versatile intermediates.
Ethanol can be used as an antiseptic to disinfect the skin before injections are given, often along with iodine. Ethanol-based soaps are becoming common in restaurants and are convenient because they do not require drying due to the volatility of the compound. Alcohol is also used as a preservative for specimens.
Alcohol gels have become common as hand sanitizers.

Production

Industrially alcohols are produced in several ways:

• By fermentation using glucose produced from sugar from the hydrolysis of starch, in the presence of yeast and temperature of less than 37 °C to produce ethanol. For instance the conversion of invertase to glucose and fructose or the conversion of glucose to zymase and ethanol.
• By direct hydration using ethylene (ethylene hydration)^[5] or other alkenes from cracking of fractions of distilled crude oil.

Endogenous

Several of the benign bacteria in the intestine use fermentation as a form of anaerobic respiration. This metabolic reaction produces ethanol as a waste product, just like aerobic respiration produces carbon dioxide and water. Thus, human bodies inevitably contain some quantity of alcohol endogenously produced by these bacteria.

Laboratory synthesis

Several methods exist for the preparation of alcohols in the laboratory.

Substitution

Primary alkyl halides react with aqueous NaOH or KOH mainly to primary alcohols in nucleophilic aliphatic substitution. (Secondary and especially tertiary alkyl halides will give the elimination (alkene) product instead). Grignard reagents react with carbonyl groups to secondary and tertiary alcohols. Related reactions are the Barbier reaction and the Nozaki-Hiyama reaction.

Reduction

Aldehydes or ketones are reduced with sodium borohydride or lithium aluminium hydride (after an acidic workup). Another reduction by aluminiumisopropylates is the Meerwein-Ponndorf-Verley reduction. Noyori asymmetric hydrogenation is the asymmetric reduction of β-keto-esters

Hydrolysis

Alkenes engage in an acid catalysed hydration reaction using concentrated sulfuric acid as a catalyst which gives usually secondary or tertiary alcohols. The hydroboration-oxidation and oxymercuration-reduction of alkenes are more reliable in organic synthesis. Alkenes react with NBS and water in halohydrin formation reaction. Amines can be converted to diazonium salts which are then hydrolyzed.
The formation of a secondary alcohol via reduction and hydration is shown:

Reactions

Deprotonation

Alcohols can behave as weak acids, undergoing deprotonation. The deprotonation reaction to produce an alkoxide salt is either performed with a strong base such as sodium hydride or n-butyllithium, or with sodium or potassium metal.

2 R-OH + 2 NaH → 2 R-O^-Na⁺ + 2H₂↑

2 R-OH + 2Na → 2R-O⁻Na + H₂

E.g. 2 CH₃CH₂-OH + 2 Na → 2 CH₃-CH₂-O⁻Na + H₂

Water is similar in pK_a to many alcohols, so with sodium hydroxide there is an equilibrium set up which usually lies to the left:

R-OH + NaOH <=> R-O^-Na⁺ + H₂O (equilibrium to the left)

It should be noted, though, that the bases used to deprotonate alcohols are strong themselves. The bases used and the alkoxides created are both highly moisture sensitive chemical reagents.
The acidity of alcohols is also affected by the overall stability of the alkoxide ion. Electron-withdrawing groups attached to the carbon containing the hydroxyl group will serve to stabilize the alkoxide when formed, thus resulting in greater acidity. On the other hand, the presence of electron-donating group will result in a less stable alkoxide ion formed. This will result in a scenario whereby the unstable alkoxide ion formed will tend to accept a proton to reform the original alcohol.
With alkyl halides alkoxides give rise to ethers in the Williamson ether synthesis.

Nucleophilic substitution

The OH group is not a good leaving group in nucleophilic substitution reactions, so neutral alcohols do not react in such reactions. However, if the oxygen is first protonated to give R−OH₂⁺, the leaving group (water) is much more stable, and the nucleophilic substitution can take place. For instance, tertiary alcohols react with hydrochloric acid to produce tertiary alkyl halides, where the hydroxyl group is replaced by a chlorine atom by unimolecular nucleophilic substitution. If primary or secondary alcohols are to be reacted with hydrochloric acid, an activator such as zinc chloride is needed. Alternatively the conversion may be performed directly using thionyl chloride.^[1]

Alcohols may likewise be converted to alkyl bromides using hydrobromic acid or phosphorus tribromide, for example:

3 R-OH + PBr₃ → 3 RBr + H₃PO₃

In the Barton-McCombie deoxygenation an alcohol is deoxygenated to an alkane with tributyltin hydride or a trimethylborane-water complex in a radical substitution reaction.

Dehydration

Alcohols are themselves nucleophilic, so R−OH₂⁺ can react with ROH to produce ethers and water in a dehydration reaction, although this reaction is rarely used except in the manufacture of diethyl ether.
More useful is the E1 elimination reaction of alcohols to produce alkenes. The reaction generally obeys Zaitsev's Rule, which states that the most stable (usually the most substituted) alkene is formed. Tertiary alcohols eliminate easily at just above room temperature, but primary alcohols require a higher temperature.
This is a diagram of acid catalysed dehydration of ethanol to produce ethene:

A more controlled elimination reaction is the Chugaev elimination with carbon disulfide and iodomethane.

Esterification

To form an ester from an alcohol and a carboxylic acid the reaction, known as Fischer esterification, is usually performed at reflux with a catalyst of concentrated sulfuric acid:

R-OH + R'-COOH → R'-COOR + H₂O

In order to drive the equilibrium to the right and produce a good yield of ester, water is usually removed, either by an excess of H₂SO₄ or by using a Dean-Stark apparatus. Esters may also be prepared by reaction of the alcohol with an acid chloride in the presence of a base such as pyridine.
Other types of ester are prepared similarly—for example tosyl (tosylate) esters are made by reaction of the alcohol with p-toluenesulfonyl chloride in pyridine.

Oxidation

Main article: Alcohol oxidation

Primary alcohols (R-CH₂-OH) can be oxidized either to aldehydes (R-CHO) or to carboxylic acids (R-CO₂H), while the oxidation of secondary alcohols (R¹R²CH-OH) normally terminates at the ketone (R¹R²C=O) stage. Tertiary alcohols (R¹R²R³C-OH) are resistant to oxidation.
The direct oxidation of primary alcohols to carboxylic acids normally proceeds via the corresponding aldehyde, which is transformed via an aldehyde hydrate (R-CH(OH)₂) by reaction with water before it can be further oxidized to the carboxylic acid.

Mechanism of oxidation of primary alcohols to carboxylic acids via aldehydes and aldehyde hydrates

Reagents useful for the transformation of primary alcohols to aldehydes are normally also suitable for the oxidation of secondary alcohols to ketones. These include Collins reagent and Dess-Martin periodinane. The direct oxidation of primary alcohols to carboxylic acids can be carried out using Potassium permanganate or the Jones reagent.

Toxicity

Main articles: Short-term effects of alcohol and Long-term effects of alcohol

Most significant of the possible long-term effects of ethanol. Additionally, in pregnant women, it causes fetal alcohol syndrome.

Ethanol in alcoholic beverages has been consumed by humans since prehistoric times for a variety of hygienic, dietary, medicinal, religious, and recreational reasons. The consumption of large doses of ethanol causes drunkenness (intoxication), which may lead to a hangover as its effects wear off. Depending upon the dose and the regularity of its consumption, ethanol can cause acute respiratory failure or death. Because ethanol impairs judgment in humans, it can be a catalyst for reckless or irresponsible behavior. The LD₅₀ of ethanol in rats is 10.3 g/kg.^[6]
Other alcohols are substantially more poisonous than ethanol, partly because they take much longer to be metabolized and partly because their metabolism produces substances that are even more toxic. Methanol (wood alcohol), for instance, is oxidized to formaldehyde and then to the poisonous formic acid in the liver by alcohol dehydrogenase and formaldehyde dehydrogenase enzymes respectively; accumulation of formic acid can lead to blindness or death.^[7] Similarly poisoning due to other alcohols such as ethylene glycol or diethylene glycol are due to their metabolites which are also produced by alcohol dehydrogenase.^[8][9] An effective treatment to prevent toxicity after methanol or ethylene glycol ingestion is to administer ethanol. Alcohol dehydrogenase has a higher affinity for ethanol, thus preventing methanol from binding and acting as a substrate. Any remaining methanol will then have time to be excreted through the kidneys.^[7][10][11]
Methanol itself, while poisonous, has a much weaker sedative effect than ethanol. Some longer-chain alcohols such as n-propanol, isopropanol, n-butanol, t-butanol and 2-methyl-2-butanol do however have stronger sedative effects, but also have higher toxicity than ethanol.^[12][13] These longer chain alcohols are found as contaminants in some alcoholic beverages and are known as fusel alcohols,^[14][15] and are reputed to cause severe hangovers although it is unclear if the fusel alcohols are actually responsible.^[16] Many longer chain alcohols are used in industry as solvents and are occasionally abused by alcoholics,^[17][18] leading to a range of adverse health effects

Electrochemistry

 Electrochemistry is a branch of chemistry that studies chemical reactions which take place in a solution at the interface of an electron conductor (a metal or a semiconductor) and an ionic conductor (the electrolyte), and which involve electron transfer between the electrode and the electrolyte or species in solution.

If a chemical reaction is driven by an external applied voltage, as in electrolysis, or if a voltage is created by a chemical reaction as in a battery, it is an electrochemical reaction. In contrast, chemical reactions where electrons are transferred between molecules are called oxidation/reduction (redox) reactions. In general, electrochemistry deals with situations where oxidation and reduction reactions are separated in space or time, connected by an external electric circuit to understand each process.

English chemists John Daniell (left) and Michael Faraday (right), both credited as founders of electrochemistry today.

 

 

 

 

 

 

16th to 18th century developments

German physicist Otto von Guericke beside his electrical generator while conducting an experiment.

Understanding of electrical matters began in the sixteenth century. During this century the English scientist William Gilbert spent 17 years experimenting with magnetism and, to a lesser extent, electricity. For his work on magnets, Gilbert became known as the "Father of Magnetism." He discovered various methods for producing and strengthening magnets.^[1]
In 1663 the German physicist Otto von Guericke created the first electric generator, which produced static electricity by applying friction in the machine. The generator was made of a large sulfur ball cast inside a glass globe, mounted on a shaft. The ball was rotated by means of a crank and a static electric spark was produced when a pad was rubbed against the ball as it rotated. The globe could be removed and used as source for experiments with electricity.^[2]
By the mid—18th century the French chemist Charles François de Cisternay du Fay had discovered two types of static electricity, and that like charges repel each other whilst unlike charges attract. Du Fay announced that electricity consisted of two fluids: "vitreous" (from the Latin for "glass"), or positive, electricity; and "resinous," or negative, electricity. This was the two-fluid theory of electricity, which was to be opposed by Benjamin Franklin's one-fluid theory later in the century.^[3]

Late 1780s diagram of Galvani's experiment on frog legs.

Charles-Augustin de Coulomb developed the law of electrostatic attraction in 1785 as an outgrowth of his attempt to investigate the law of electrical repulsions as stated by Joseph Priestley in England.^[4]

Italian physicist Alessandro Volta showing his "battery" to French emperor Napoleon Bonaparte in the early 19th century.

In the late 18th century the Italian physician and anatomist Luigi Galvani marked the birth of electrochemistry by establishing a bridge between chemical reactions and electricity on his essay "De Viribus Electricitatis in Motu Musculari Commentarius" (Latin for Commentary on the Effect of Electricity on Muscular Motion) in 1791 where he proposed a "nerveo-electrical substance" on biological life forms.^[5]
In his essay Galvani concluded that animal tissue contained a here-to-fore neglected innate, vital force, which he termed "animal electricity," which activated nerves and muscles spanned by metal probes. He believed that this new force was a form of electricity in addition to the "natural" form produced by lightning or by the electric eel and torpedo ray as well as the "artificial" form produced by friction (i.e., static electricity).^[6]
Galvani's scientific colleagues generally accepted his views, but Alessandro Volta rejected the idea of an "animal electric fluid," replying that the frog's legs responded to differences in metal temper, composition, and bulk.^[5][6] Galvani refuted this by obtaining muscular action with two pieces of the same material.

[edit] 19th century

Sir Humphry Davy's portrait in the 19th century.

In 1800, William Nicholson and Johann Wilhelm Ritter succeeded in decomposing water into hydrogen and oxygen by electrolysis. Soon thereafter Ritter discovered the process of electroplating. He also observed that the amount of metal deposited and the amount of oxygen produced during an electrolytic process depended on the distance between the electrodes.^[7] By 1801 Ritter observed thermoelectric currents and anticipated the discovery of thermoelectricity by Thomas Johann Seebeck.^[8]
By the 1810s William Hyde Wollaston made improvements to the galvanic cell. Sir Humphry Davy's work with electrolysis led to the conclusion that the production of electricity in simple electrolytic cells resulted from chemical action and that chemical combination occurred between substances of opposite charge. This work led directly to the isolation of sodium and potassium from their compounds and of the alkaline earth metals from theirs in 1808.^[9]
Hans Christian Ørsted's discovery of the magnetic effect of electrical currents in 1820 was immediately recognized as an epoch-making advance, although he left further work on electromagnetism to others. André-Marie Ampère quickly repeated Oestred's experiment, and formulated them mathematically.^[10]
In 1821, Estonian-German physicist Thomas Johann Seebeck demonstrated the electrical potential in the juncture points of two dissimilar metals when there is a heat difference between the joints.^[11]
In 1827, the German scientist Georg Ohm expressed his law in this famous book "Die galvanische Kette, mathematisch bearbeitet" (The Galvanic Circuit Investigated Mathematically) in which he gave his complete theory of electricity.^[11]
In 1832, Michael Faraday's experiments led him to state his two laws of electrochemistry. In 1836, John Daniell invented a primary cell in which hydrogen was eliminated in the generation of the electricity. Daniell had solved the problem of polarization. Later results revealed that alloying the amalgamated zinc with mercury would produce a better voltage.

Swedish chemist Svante Arrhenius portrait circa 1880s.

William Grove produced the first fuel cell in 1839. In 1846, Wilhelm Weber developed the electrodynamometer. In 1868, Georges Leclanché patented a new cell which eventually became the forerunner to the world's first widely used battery, the zinc carbon cell.^[7]
Svante Arrhenius published his thesis in 1884 on Recherches sur la conductibilité galvanique des électrolytes (Investigations on the galvanic conductivity of electrolytes). From his results the author concluded that electrolytes, when dissolved in water, become to varying degrees split or dissociated into electrically opposite positive and negative ions.^[12]
In 1886, Paul Héroult and Charles M. Hall developed an efficient method (the Hall–Héroult process) to obtain aluminium using elecrolysis of molten alumina.^[13]
In 1894, Friedrich Ostwald concluded important studies of the conductivity and electrolytic dissociation of organic acids.^[14]

German scientist Walther Nernst portrait in the 1910s.

Walther Hermann Nernst developed the theory of the electromotive force of the voltaic cell in 1888. In 1889, he showed how the characteristics of the current produced could be used to calculate the free energy change in the chemical reaction producing the current. He constructed an equation, known as Nernst equation, which related the voltage of a cell to its properties.^[15]
In 1898, Fritz Haber showed that definite reduction products can result from electrolytic processes if the potential at the cathode is kept constant. In 1898, he explained the reduction of nitrobenzene in stages at the cathode and this became the model for other similar reduction processes.^[16]

[edit] The 20th century and recent developments

In 1902, The Electrochemical Society (ECS) was founded.^[17]
In 1909, Robert Andrews Millikan began a series of experiments to determine the electric charge carried by a single electron.^[18]
In 1923, Johannes Nicolaus Brønsted and Martin Lowry published essentially the same theory about how acids and bases behave, using an electrochemical basis.^[19]
Arne Tiselius developed the first sophisticated electrophoretic apparatus in 1937 and some years later he was awarded the 1948 Nobel Prize for his work in protein electrophoresis.^[20]
A year later, in 1949, the International Society of Electrochemistry (ISE) was founded.^[21]
By the 1960s–1970s quantum electrochemistry was developed by Revaz Dogonadze and his pupils.

[edit] Principles

[edit] Redox reactions

Main article: Redox reaction

Redox stands for reduction-oxidation, and are electrochemical processes involving electron transfer to or from a molecule or ion changing its oxidation state. This reaction can occur through the application of an external voltage or through the release of chemical energy.

[edit] Oxidation and reduction

Oxidation and reduction describe the change of oxidation state that takes place in the atoms, ions or molecules involved in an electrochemical reaction. Formally, oxidation state is the hypothetical charge that an atom would have if all bonds to atoms of different elements were 100% ionic. An atom or ion that gives up an electron to another atom or ion has its oxidation state increase, and the recipient of the negatively charged electron has its oxidation state decrease. Oxidation and reduction always occur in a paired fashion such that one species is oxidized when another is reduced. This paired electron transfer is called a redox reaction.
For example, when atomic sodium reacts with atomic chlorine, sodium donates one electron and attains an oxidation state of +1. Chlorine accepts the electron and its oxidation state is reduced to −1. The sign of the oxidation state (positive/negative) actually corresponds to the value of each ion's electronic charge. The attraction of the differently charged sodium and chlorine ions is the reason they then form an ionic bond.
The loss of electrons from an atom or molecule is called oxidation, and the gain of electrons is reduction. This can be easily remembered through the use of mnemonic devices. Two of the most popular are "OIL RIG" (Oxidation Is Loss, Reduction Is Gain) and "LEO" the lion says "GER" (Lose Electrons: Oxidization, Gain Electrons: Reduction). For cases where electrons are shared (covalent bonds) between atoms with large differences in electronegativity, the electron is assigned to the atom with the largest electronegativity in determining the oxidation state.
The atom or molecule which loses electrons is known as the reducing agent, or reductant, and the substance which accepts the electrons is called the oxidizing agent, or oxidant. The oxidizing agent is always being reduced in a reaction; the reducing agent is always being oxidized. Oxygen is a common oxidizing agent, but not the only one. Despite the name, an oxidation reaction does not necessarily need to involve oxygen. In fact, a fire can be fed by an oxidant other than oxygen; fluorine fires are often unquenchable, as fluorine is an even stronger oxidant (it has a higher electronegativity) than oxygen.
For reactions involving oxygen, the gain of oxygen implies the oxidation of the atom or molecule to which the oxygen is added (and the oxygen is reduced). In organic compounds, such as butane or ethanol, the loss of hydrogen implies oxidation of the molecule from which it is lost (and the hydrogen is reduced). This follows because the hydrogen donates its electron in covalent bonds with non-metals but it takes the electron along when it is lost. Conversely, loss of oxygen or gain of hydrogen implies reduction.

[edit] Balancing redox reactions

Main article: Chemical equation

Electrochemical reactions in water are better understood by balancing redox reactions using the ion-electron method where H⁺ , OH^– ion, H₂O and electrons (to compensate the oxidation changes) are added to cell's half-reactions for oxidation and reduction.

[edit] Acidic medium

In acid medium H⁺ ions and water are added to half-reactions to balance the overall reaction. For example, when manganese reacts with sodium bismuthate.

Unbalanced reaction: Mn²⁺(aq) + NaBiO₃(s) → Bi³⁺(aq) + MnO₄^–(aq)

Oxidation: 4 H₂O(l) + Mn²⁺(aq) → MnO₄^–(aq) + 8 H⁺(aq) + 5 e^–

Reduction: 2 e^– + 6 H⁺(aq) + BiO₃^–(s) → Bi³⁺(aq) + 3 H₂O(l)

Finally, the reaction is balanced by multiplying the number of electrons from the reduction half reaction to oxidation half reaction and vice versa and adding both half reactions, thus solving the equation.

8 H₂O(l) + 2 Mn²⁺(aq) → 2 MnO₄^–(aq) + 16 H⁺(aq) + 10 e^–

10 e^– + 30 H⁺(aq) + 5 BiO₃^–(s) → 5 Bi³⁺(aq) + 15 H₂O(l)

Reaction balanced:

14 H⁺(aq) + 2 Mn²⁺(aq) + 5 NaBiO₃(s) → 7 H₂O(l) + 2 MnO₄^–(aq) + 5 Bi³⁺(aq) + 5 Na⁺(aq)

[edit] Basic medium

In basic medium OH^– ions and water are added to half reactions to balance the overall reaction. For example, on reaction between potassium permanganate and sodium sulfite.

Unbalanced reaction: KMnO₄ + Na₂SO₃ + H₂O → MnO₂ + Na₂SO₄ + KOH

Reduction: 3 e^– + 2 H₂O + MnO₄^– → MnO₂ + 4 OH^–

Oxidation: 2 OH^– + SO₃^2– → SO₄^2– + H₂O + 2 e^–

The same procedure as followed on acid medium by multiplying electrons to opposite half reactions solve the equation thus balancing the overall reaction.

6 e^– + 4 H₂O + 2 MnO₄^– → 2 MnO₂ + 8 OH^–

6 OH^– + 3 SO₃^2– → 3 SO₄^2– + 3 H₂O + 6e^–

Equation balanced:

2 KMnO₄ + 3 Na₂SO₃ + H₂O → 2 MnO₂ + 3 Na₂SO₄ + 2 KOH

[edit] Neutral medium

The same procedure as used on acid medium is applied, for example on balancing using electron ion method to complete combustion of propane.

Unbalanced reaction: C₃H₈ + O₂ → CO₂ + H₂O

Reduction: 4 H⁺ + O₂ + 4 e^– → 2 H₂O

Oxidation: 6 H₂O + C₃H₈ → 3 CO₂ + 20 e^– + 20 H⁺

As in acid and basic medium, electrons which were used to compensate oxidation changes are multiplied to opposite half reactions, thus solving the equation.

20 H⁺ + 5 O₂ + 20 e^– → 10 H₂O

6 H₂O + C₃H₈ → 3 CO₂ + 20 e^– + 20 H⁺

Equation balanced:

C₃H₈ + 5 O₂ → 3 CO₂ + 4 H₂O

[edit] Electrochemical cells

Main article: Electrochemical cell

An electrochemical cell is a device that produces an electric current from energy released by a spontaneous redox reaction. This kind of cell includes the Galvanic cell or Voltaic cell, named after Luigi Galvani and Alessandro Volta, both scientists who conducted several experiments on chemical reactions and electric current during the late 18th century.
Electrochemical cells have two conductive electrodes (the anode and the cathode). The anode is defined as the electrode where oxidation occurs and the cathode is the electrode where the reduction takes place. Electrodes can be made from any sufficiently conductive materials, such as metals, semiconductors, graphite, and even conductive polymers. In between these electrodes is the electrolyte, which contains ions that can freely move.
The Galvanic cell uses two different metal electrodes, each in an electrolyte where the positively charged ions are the oxidized form of the electrode metal. One electrode will undergo oxidation (the anode) and the other will undergo reduction (the cathode). The metal of the anode will oxidize, going from an oxidation state of 0 (in the solid form) to a positive oxidation state and become an ion. At the cathode, the metal ion in solution will accept one or more electrons from the cathode and the ion's oxidation state is reduced to 0. This forms a solid metal that electrodeposits on the cathode. The two electrodes must be electrically connected to each other, allowing for a flow of electrons that leave the metal of the anode and flow through this connection to the ions at the surface of the cathode. This flow of electrons is an electrical current that can be used to do work, such as turn a motor or power a light.
A Galvanic cell whose electrodes are zinc and copper submerged in zinc sulfate and copper sulfate, respectively, is known as a Daniell cell.^[22]
Half reactions for a Daniell cell are these:^[22]

Zinc electrode (anode): Zn(s) → Zn²⁺(aq) + 2 e^–

Copper electrode (cathode): Cu²⁺(aq) + 2 e^– → Cu(s)

A modern cell stand for electrochemical research. The electrodes attach to high-quality metallic wires, and the stand is attached to a potentiostat/galvanostat (not pictured). A shot glass-shaped container is aerated with a noble gas and sealed with the Teflon block.

In this example, the anode is zinc metal which oxidizes (loses electrons) to form zinc ions in solution, and copper ions accept electrons from the copper metal electrode and the ions deposit at the copper cathode as an electrodeposit. This cell forms a simple battery as it will spontaneously generate a flow of electrical current from the anode to the cathode through the external connection. This reaction can be driven in reverse by applying a voltage, resulting in the deposition of zinc metal at the anode and formation of copper ions at the cathode.^[22]
To provide a complete electric circuit, there must also be an ionic conduction path between the anode and cathode electrolytes in addition to the electron conduction path. The simplest ionic conduction path is to provide a liquid junction. To avoid mixing between the two electrolytes, the liquid junction can be provided through a porous plug that allows ion flow while reducing electrolyte mixing. To further minimize mixing of the electrolytes, a salt bridge can be used which consists of an electrolyte saturated gel in an inverted U-tube. As the negatively charged electrons flow in one direction around this circuit, the positively charged metal ions flow in the opposite direction in the electrolyte.
A voltmeter is capable of measuring the change of electrical potential between the anode and the cathode.
Electrochemical cell voltage is also referred to as electromotive force or emf.
A cell diagram can be used to trace the path of the electrons in the electrochemical cell. For example, here is a cell diagram of a Daniell cell:

Zn(s) | Zn²⁺ (1M) || Cu²⁺ (1M) | Cu(s)

First, the reduced form of the metal to be oxidized at the anode (Zn) is written. This is separated from its oxidized form by a vertical line, which represents the limit between the phases (oxidation changes). The double vertical lines represent the saline bridge on the cell. Finally, the oxidized form of the metal to be reduced at the cathode, is written, separated from its reduced form by the vertical line. The electrolyte concentration is given as it is an important variable in determining the cell potential.

[edit] Standard electrode potential

Main article: Standard electrode potential

To allow prediction of the cell potential, tabulations of standard electrode potential are available. Such tabulations are referenced to the standard hydrogen electrode (SHE). The standard hydrogen electrode undergoes the reaction

2 H⁺(aq) + 2 e^– → H₂

which is shown as reduction but, in fact, the SHE can act as either the anode or the cathode, depending on the relative oxidation/reduction potential of the other electrode/electrolyte combination. The term standard in SHE requires a supply of hydrogen gas bubbled through the electrolyte at a pressure of 1 atm and an acidic electrolyte with H⁺ activity equal to 1 (usually assumed to be [H⁺] = 1 mol/liter).
The SHE electrode can be connected to any other electrode by a salt bridge to form a cell. If the second electrode is also at standard conditions, then the measured cell potential is called the standard electrode potential for the electrode. The standard electrode potential for the SHE is zero, by definition. The polarity of the standard electrode potential provides information about the relative reduction potential of the electrode compared to the SHE. If the electrode has a positive potential with respect to the SHE, then that means it is a strongly reducing electrode which forces the SHE to be the anode (an example is Cu in aqueous CuSO₄ with a standard electrode potential of 0.337 V). Conversely, if the measured potential is negative, the electrode is more oxidizing than the SHE (such as Zn in ZnSO₄ where the standard electrode potential is −0.76 V).^[22]
Standard electrode potentials are usually tabulated as reduction potentials. However, the reactions are reversible and the role of a particular electrode in a cell depends on the relative oxidation/reduction potential of both electrodes. The oxidation potential for a particular electrode is just the negative of the reduction potential. A standard cell potential can be determined by looking up the standard electrode potentials for both electrodes (sometimes called half cell potentials). The one that is smaller will be the anode and will undergo oxidation. The cell potential is then calculated as the sum of the reduction potential for the cathode and the oxidation potential for the anode.

E°_cell = E°_red(cathode) – E°_red(anode) = E°_red(cathode) + E°_oxi(anode)

For example, the standard electrode potential for a copper electrode is:

Cell diagram

Pt(s) | H₂(1 atm) | H⁺(1 M) || Cu²⁺ (1 M) | Cu(s)

E°_cell = E°_red(cathode) – E°_red(anode)

At standard temperature, pressure and concentration conditions, the cell's emf (measured by a multimeter) is 0.34 V. By definition, the electrode potential for the SHE is zero. Thus, the Cu is the cathode and the SHE is the anode giving

E_cell = E°(Cu²⁺/Cu) – E°(H⁺/H₂)

Or,

E°(Cu²⁺/Cu) = 0.34 V

Changes in the stoichiometric coefficients of a balanced cell equation will not change E°_red value because the standard electrode potential is an intensive property.

[edit] Spontaneity of redox reaction

Main article: Spontaneous process

During operation of electrochemical cells, chemical energy is transformed into electrical energy and is expressed mathematically as the product of the cell's emf and the electrical charge transferred through the external circuit.

Electrical energy = E_cellC_trans

where E_cell is the cell potential measured in volts (V) and C_trans is the cell current integrated over time and measured in coulombs (C); C_trans can also be determined by multiplying the total number of electrons transferred (measured in moles) times Faraday's constant (F).
The emf of the cell at zero current is the maximum possible emf. It is used to calculate the maximum possible electrical energy that could be obtained from a chemical reaction. This energy is referred to as electrical work and is expressed by the following equation:

W_max = W_electrical = –nF·E_cell

where work is defined as positive into the system.
Since the free energy is the maximum amount of work that can be extracted from a system, one can write:^[23]

ΔG = –nF·E_cell

A positive cell potential gives a negative change in Gibbs free energy. This is consistent with the cell production of an electric current flowing from the cathode to the anode through the external circuit. If the current is driven in the opposite direction by imposing an external potential, then work is done on the cell to drive electrolysis.^[23]
A spontaneous electrochemical reaction (change in Gibbs free energy less than zero) can be used to generate an electric current, in electrochemical cells. This is the basis of all batteries and fuel cells. For example, gaseous oxygen (O₂) and hydrogen (H₂) can be combined in a fuel cell to form water and energy, typically a combination of heat and electrical energy.^[23]
Conversely, non-spontaneous electrochemical reactions can be driven forward by the application of a current at sufficient voltage. The electrolysis of water into gaseous oxygen and hydrogen is a typical example.
The relation between the equilibrium constant, K, and the Gibbs free energy for an electrochemical cell is expressed as follows:

ΔG° = –RT ln(K) = –nF·E°_cell

Rearranging to express the relation between standard potential and equilibrium constant yields

Previous equation can use Briggsian logarithm as shown below:

[edit] Cell emf dependency on changes in concentration

[edit] Nernst equation

Main article: Nernst equation

The standard potential of an electrochemical cell requires standard conditions for all of the reactants. When reactant concentrations differ from standard conditions, the cell potential will deviate from the standard potential. In the 20th century German chemist Walther Nernst proposed a mathematical model to determine the effect of reactant concentration on electrochemical cell potential.
In the late 19th century, Josiah Willard Gibbs had formulated a theory to predict whether a chemical reaction is spontaneous based on the free energy

ΔG = ΔG° + RT·ln(Q)

Here ΔG is change in Gibbs free energy, T is absolute temperature, R is the gas constant and Q is reaction quotient.
Gibbs' key contribution was to formalize the understanding of the effect of reactant concentration on spontaneity.
Based on Gibbs' work, Nernst extended the theory to include the contribution from electric potential on charged species. As shown in the previous section, the change in Gibbs free energy for an electrochemical cell can be related to the cell potential. Thus, Gibbs' theory becomes

nFΔE = nFΔE° – RT ln(Q)

Here n is the number of electrons/mole product, F is the Faraday constant (coulombs/mole), and ΔE is cell potential.
Finally, Nernst divided through by the amount of charge transferred to arrive at a new equation which now bears his name:

ΔE = ΔE° – (RT/nF)ln(Q)

Assuming standard conditions (T = 25 °C) and R = 8.3145 J/(K·mol), the equation above can be expressed on base—10 logarithm as shown below:^[24]

[edit] Concentration cells

Main article: Concentration cell

A concentration cell is an electrochemical cell where the two electrodes are the same material, the electrolytes on the two half-cells involve the same ions, but the electrolyte concentration differs between the two half-cells.
For example an electrochemical cell, where two copper electrodes are submerged in two copper(II) sulfate solutions, whose concentrations are 0.05 M and 2.0 M, connected through a salt bridge. This type of cell will generate a potential that can be predicted by the Nernst equation. Both electrodes undergo the same chemistry (although the reaction proceeds in reverse at the cathode)

Cu²⁺(aq) + 2 e^– → Cu(s)

Le Chatelier's principle indicates that the reaction is more favorable to reduction as the concentration of Cu²⁺ ions increases. Reduction will take place in the cell's compartment where concentration is higher and oxidation will occur on the more dilute side.
The following cell diagram describes the cell mentioned above:

Cu(s) | Cu²⁺ (0.05 M) || Cu²⁺ (2.0 M) | Cu(s)

Where the half cell reactions for oxidation and reduction are:

Oxidation: Cu(s) → Cu²⁺ (0.05 M) + 2 e^–

Reduction: Cu²⁺ (2.0 M) + 2 e^– → Cu(s)

Overall reaction: Cu²⁺ (2.0 M) → Cu²⁺ (0.05 M)

The cell's emf is calculated through Nernst equation as follows:

The value of E° in this kind of cell is zero, as electrodes and ions are the same in both half-cells.
After replacing values from the case mentioned, it is possible to calculate cell's potential:

or by:

However, this value is only approximate, as reaction quotient is defined in terms of ion activities which can be approximated with the concentrations as calculated here.
The Nernst equation plays an important role in understanding electrical effects in cells and organelles. Such effects include nerve synapses and cardiac beat as well as the resting potential of a somatic cell.

[edit] Battery

Main article: Battery (electricity)

A battery is a number of cells combined and is used to supply electrical energy that is stored chemically. The term battery is often, and incorrectly, used to describe a single cell. In a battery, cells are usually wired in series to increase the supply voltage but sometimes wired in parallel to allow greater current to be supplied. Batteries are optimized to produce a constant electric current for as long as possible. Although the cells discussed previously are useful for theoretical purposes and some laboratory experiments, the large internal resistance of the salt bridge make them inappropriate battery technologies. Various alternative battery technologies have been commercialized as discussed next.

[edit] Dry cell

Main article: Dry cell

Zinc carbon battery diagram.

Dry cells do not have a fluid electrolyte. Instead, they use a moist electrolyte paste. Leclanché's cell is a good example of this, where the anode is a zinc container surrounded by a thin layer of manganese dioxide and a moist electrolyte paste of ammonium chloride and zinc chloride mixed with starch. The cell's cathode is represented by a carbon bar inserted on the cell's electrolyte, usually placed in the middle.^[25]
Leclanché's simplified half reactions are shown below:^[25]

Anode: Zn(s) → Zn²⁺(aq) + 2 e^–

Cathode: 2 NH₄⁺(aq) + 2 MnO₂(s) + 2 e^– → Mn₂O₃(s) + 2 NH₃(aq) + H₂O(l)

Overall reaction: Zn(s) + 2 NH₄⁺(aq) + 2 MnO₂(s) → Zn²⁺(aq) + Mn₂O₃(s) + 2 NH₃(aq) + H₂O(l)

The voltage obtained from the zinc-carbon battery is around 1.5 V.^[25]

[edit] Mercury battery

Main article: Mercury battery

Cutaway view of a mercury battery.

This battery first appeared in the early 1940s. The mercury battery has many applications in medicine and electronics. The battery consists of a steel—made container in the shape of a cylinder acting as the cathode, where an amalgamated anode of mercury and zinc is surrounded by a stronger alkaline electrolyte and a paste of zinc oxide and mercury(II) oxide.^[26]
Mercury battery half reactions are shown below:

Anode: Zn(Hg) + 2 OH^–(aq) → ZnO(s) + H₂O(l) + 2 e^–

Cathode: HgO(s) + H₂O(l) + 2 e^– → Hg(l) + 2 OH^–(aq)

Overall reaction: Zn(Hg) + HgO(s) → ZnO(s) + Hg(l)

There are no changes in the electrolyte's composition when the cell works. Such batteries provide 1.35 V of direct current.^[27]

[edit] Lead-acid battery

A sealed lead-acid battery.

Main article: Lead-acid battery

The lead-acid battery used in automobiles, consists of a series of six identical cells assembled in series. Each cell has a lead anode and a cathode made from lead dioxide packed in a metal plaque. Cathode and anode are submerged in a solution of sulfuric acid acting as the electrolyte.^[28]
Lead-acid battery half cell reactions are shown below:^[28]

Anode: Pb(s) + SO₄^2–(aq) → PbSO₄(s) + 2 e^–

Cathode: PbO₂(s) + 4 H⁺(aq) + SO₄^2–(aq) + 2 e^– → PbSO₄(s) + 2 H₂O(l)

Overall reaction: Pb(s) + PbO₂(s) + 4 H⁺(aq) + 2 SO₄^2–(aq) → 2 PbSO₄(s) + 2 H₂O(l)

At standard conditions, each cell may produce a potential of 2 V, hence overall voltage produced is 12 V. Differing from mercury and zinc-carbon batteries, lead-acid batteries are rechargeable. If an external voltage is supplied to the battery it will produce an electrolysis of the products in the overall reaction (discharge), thus recovering initial components which made the battery work.^[28]

[edit] Lithium rechargeable battery

Main article: Lithium battery

Instead of an aqueous electrolyte or a moist electrolyte paste, a solid state battery operates using a solid electrolyte. Lithium polymer batteries are an example of this; a graphite bar acts as the anode, a bar of lithium cobaltate acts as the cathode, and a polymer, swollen with a lithium salt, allows the passage of ions and serves as the electrolyte. In this cell, the carbon in the anode can reversibly form a lithium-carbon alloy. Upon discharging, lithium ions spontaneously leave the lithium cobaltate cathode and travel through the polymer and into the carbon anode forming the alloy. This flow of positive lithium ions is the electrical current that the battery provides. By charging the cell, the lithium dealloys and travels back into the cathode. The advantage of this kind of battery is that Lithium possess the highest negative value of standard reduction potential. It is also a light metal and therefore less mass is required to generate 1 mole of electrons. Lithium ion battery technologies are widely used in portable electronic devices because they have high energy storage density and are rechargeable. These technologies show promise for future automotive applications, with new materials such as iron phosphates and lithium vanadates.

Flow battery