|Acids and bases|
An acid–base reaction is a chemical reaction that occurs between an acid and a base. Several concepts exist that provide alternative definitions for the reaction mechanisms involved and their application in solving related problems. Despite several differences in definitions, their importance becomes apparent as different methods of analysis when applied to acid–base reactions for gaseous or liquid species, or when acid or base character may be somewhat less apparent. The first of these scientific concepts of acids and bases was provided by the French chemist Antoine Lavoisier, circa 1776.
Historic acid–base theories 
Lavoisier's oxygen theory of acids 
The first scientific concept of acids and bases was provided by Lavoisier circa 1776. Since Lavoisier's knowledge of strong acids was mainly restricted to oxoacids, such as HNO3 (nitric acid) and H2SO4 (sulphuric acid), which tend to contain central atoms in high oxidation states surrounded by oxygen, and since he was not aware of the true composition of the hydrohalic acids (HF, HCl, HBr, and HI), he defined acids in terms of their containing oxygen, which in fact he named from Greek words meaning "acid-former" (from the Greek οξυς (oxys) meaning "acid" or "sharp" and γεινομαι (geinomai) meaning "engender"). The Lavoisier definition was held as absolute truth for over 30 years, until the 1810 article and subsequent lectures by Sir Humphry Davy in which he proved the lack of oxygen in H2S, H2Te, and the hydrohalic acids. However, Davy failed to develop a new theory, concluding that "acidity does not depend upon any particular elementary substance, but upon peculiar arrangement of various substances". One notable modification of oxygen theory was provided by Berzelius, who stated that acids are oxides of nonmetals while bases are oxides of metals.
Liebig's hydrogen theory of acids 
Circa 1838 Justus von Liebig proposed that an acid is a hydrogen-containing substance in which the hydrogen could be replaced by a metal. This redefinition was based on his extensive work on the chemical composition of organic acids, finishing the doctrinal shift from oxygen-based acids to hydrogen-based acids started by Davy. Liebig's definition, while completely empirical, remained in use for almost 50 years until the adoption of the Arrhenius definition.
Common acid–base theories 
Arrhenius definition 
The first modern definition of acids and bases was devised by Svante Arrhenius. A hydrogen theory of acids, it followed from his 1884 work with Friedrich Wilhelm Ostwald in establishing the presence of ions in aqueous solution and led to Arrhenius receiving the Nobel Prize in Chemistry in 1903.
As defined by Arrhenius:
- an Arrhenius acid is a substance that dissociates in water to form hydrogen ions (H+); that is, an acid increases the concentration of H+ ions in an aqueous solution. This protonation of water yields hydronium (H3O+); in modern times, the use of H+ is regarded as a shorthand for H3O+, because it is now known that a bare proton (H+) does not exist as a free species in aqueous solution.
- an Arrhenius base is a substance that dissociates in water to form hydroxide (OH−) ions; that is, a base increases the concentration of OH− ions in an aqueous solution.
The Arrhenius definitions of acidity and alkalinity are restricted to aqueous solutions, and refer to the concentration of the solvent ions. Under this definition, pure H2SO4 and HCl dissolved in toluene are not acidic, and molten KOH and solutions of sodium amide in liquid ammonia are not alkaline.
The universal aqueous acid–base definition of the Arrhenius concept is described as the formation of a water molecule from a proton and hydroxide ion. This leads to the definition that in Arrhenius acid–base reactions, a salt and water are formed from the reaction between an acid and a base. This is a neutralization reaction which has been put into a word equation:
- acid + base → salt + water
The positive ion from a base and the negative ion from an acid form a salt together. For example, two moles of sodium ion (Na+) from the base sodium hydroxide (NaOH) combine with one mole of sulfate ion (SO2−
4) from sulfuric acid (H2SO4) to form one mole of sodium sulfate (Na2SO4), along with two moles of water:
- 2 NaOH + H2SO4 → Na2SO4 + 2 H2O
Solvent system definition 
One of the limitations of the Arrhenius definition is its reliance on water solutions. Edward C. Franklin studied the acid–base reactions in liquid ammonia in 1905 and pointed out the similarities to the water-based Arrhenius theory. Albert F. O. Germann, working with liquid COCl2, formulated the solvent-based theory in 1925, thereby generalizing the Arrhenius definition to cover aprotic solvents.
Germann pointed out that in many solutions, there are ions in equilibrium with the neutral solvent molecules:
- 2 H2O H3O+ + OH−
- 2 NH3 NH+
4 + NH−
Some aprotic systems also undergo such dissociation, such as dinitrogen tetroxide into nitrosonium and nitrate, antimony trichloride into dichloroantimonium and tetrachloroantimonate, and phosgene into chlorocarboxonium and chloride:
- N2O4 NO+ + NO−
- 2 SbCl3 SbCl+
2 + SbCl−
- COCl2 COCl+ + Cl−
A solute that causes an increase in the concentration of the solvonium ions and a decrease in the concentration of solvate ions is defined as an acid. A solute that causes an increase in the concentration of the solvate ions and a decrease in the concentration of the solvonium ions is defined as a base.
Thus, in liquid ammonia, KNH2 (supplying NH−
2) is a strong base, and NH4NO3 (supplying NH+
4) is a strong acid. In liquid sulfur dioxide (SO2), thionyl compounds (supplying SO2+) behave as acids, and sulfites (supplying SO2−
3) behave as bases.
The non-aqueous acid–base reactions in liquid ammonia are similar to the reactions in water:
- 2 NaNH2 (base) + Zn(NH2)2 (amphiphilic amide) → Na2[Zn(NH2)4]
- 2 NH4I (acid) + Zn(NH2)2 (amphiphilic amide) → [Zn(NH3)4)]I2
Nitric acid can be a base in liquid sulfuric acid:
- HNO3 (base) + 2 H2SO4 → NO+
2 + H3O+ + 2 HSO−
The unique strength of this definition shows in describing the reactions in aprotic solvents; for example, in liquid N2O4:
- AgNO3 (base) + NOCl (acid) → N2O4 (solvent) + AgCl (salt)
Because the solvent system definition depends on the solute as well as on the solvent itself, a particular solute can be either an acid or a base depending on the choice of the solvent: HClO4 is a strong acid in water, a weak acid in acetic acid, and a weak base in fluorosulfonic acid; this characteristic of the theory has been seen as both a strength and a weakness, because some substances (such as SO3 and NH3) have been seen to be acidic or basic on their own right. On the other hand, solvent system theory has been criticized as being too general to be useful. Also, it has been thought that there is something intrinsically acidic about hydrogen compounds, a property not shared by non-hydrogenic solvonium salts.
Brønsted–Lowry definition 
The Brønsted–Lowry definition, formulated in 1923, independently by Johannes Nicolaus Brønsted in Denmark and Martin Lowry in England, is based upon the idea of protonation of bases through the de-protonation of acids – that is, the ability of acids to "donate" hydrogen ions (H+) or protons to bases, which "accept" them. Unlike the previous definitions, the Brønsted–Lowry definition does not refer to the formation of salt and solvent, but instead to the formation of conjugate acids and conjugate bases, produced by the transfer of a proton from the acid to the base. In this approach, acids and bases are fundamentally different in behavior from salts, which are seen as electrolytes, subject to the theories of Debye, Onsager, and others. An acid and a base react not to produce a salt and a solvent, but to form a new acid and a new base. The concept of neutralization is thus absent.
According to Brønsted–Lowry definition, an acid is a compound that can donate a proton, and a base is a compound that can receive a proton. An acid–base reaction is, thus, the removal of a hydrogen ion from the acid and its addition to the base. This does not refer to the removal of a proton from the nucleus of an atom, which would require levels of energy not attainable through the simple dissociation of acids, but to removal of a hydrogen ion (H+).
The removal of a proton (hydrogen ion) from an acid produces its conjugate base, which is the acid with a hydrogen ion removed, and the reception of a proton by a base produces its conjugate acid, which is the base with a hydrogen ion added.
For example, the removal of H+ from hydrochloric acid (HCl) produces the chloride ion (Cl−), the conjugate base of the acid:
- HCl → H+ + Cl−
The addition of H+ to the hydroxide ion (OH−), a base, produces water (H2O), its conjugate acid:
- H+ + OH− → H2O
Although Brønsted–Lowry acid–base behavior is formally independent of any solvent, it encompasses Arrhenius and solvent system definitions in an unenforced way. For example, protonation of ammonia, a base, gives ammonium ion, its conjugate acid:
- H+ + NH3 → NH+
The reaction of ammonia, a base, with acetic acid in absence of water can be described to give ammonium cation, an acid, and acetate anion, a base:
- CH3COOH + NH3 → NH+
4 + CH3COO−
This definition also explains the dissociation of water into low concentrations of hydronium and hydroxide ions:
- H2O + H2O H3O+ + OH−
Water, being amphoteric, can act as both an acid and a base; here, one molecule of water acts as an acid, donating a H+ ion and forming the conjugate base, OH−, and a second molecule of water acts as a base, accepting the H+ ion and forming the conjugate acid, H3O+.
- HCl (acid) + H2O (base) H3O+ (acid) + Cl− (base)
4 (acid) + H2O (base) H3O+ (acid) + NH3 (base)
as are basic dissociation and basic hydrolysis:
- NH3 (base) + H2O (acid) NH+
4 (acid) + OH− (base)
- CH3COO− (base) + H2O (acid) CH3COOH (acid) + OH− (base)
Thus, the general formula for acid–base reactions according to the Brønsted–Lowry definition is:
- AH + B → BH+ + A−
where AH represents the acid, B represents the base, BH+ represents the conjugate acid of B, and A− represents the conjugate base of AH.
Although Brønsted–Lowry calls hydrogen-containing substances like HCl acids, KOH and KNH2 are not bases but salts containing the bases OH− and NH−
2. Also, some substances, which many chemists considered to be acids, such as SO3 or BCl3, are excluded from this classification due to lack of hydrogen. Gilbert Lewis wrote in 1938, "To restrict the group of acids to those substances that contain hydrogen interferes as seriously with the systematic understanding of chemistry as would the restriction of the term oxidizing agent to substances containing oxygen."
Lewis definition 
The hydrogen requirement of Arrhenius and Brønsted–Lowry was removed by the Lewis definition of acid–base reactions, devised by Gilbert N. Lewis in 1923, in the same year as Brønsted–Lowry, but it was not elaborated by him until 1938. Instead of defining acid–base reactions in terms of protons or other bonded substances, the Lewis definition defines a base (referred to as a Lewis base) to be a compound that can donate an electron pair, and an acid (a Lewis acid) to be a compound that can receive this electron pair.
In this system, an acid does not exchange atoms with a base, but combines with it. For example, consider this classical aqueous acid–base reaction:
- HCl (aq) + NaOH (aq) → H2O (l) + NaCl (aq)
The Lewis definition does not regard this reaction as the formation of salt and water or the transfer of H+ from HCl to OH−. Instead, it regards the acid to be the H+ ion itself, and the base to be the OH− ion, which has an unshared electron pair. Therefore, the acid–base reaction here, according to the Lewis definition, is the donation of the electron pair from OH− to the H+ ion. This forms a covalent bond between H+ and OH−, thus producing water (H2O).
By treating acid–base reactions in terms of electron pairs instead of specific substances, the Lewis definition can be applied to reactions that do not fall under other definitions of acid–base reactions. For example, a silver cation behaves as an acid with respect to ammonia, which behaves as a base, in the following reaction:
- Ag+ + 2 :NH3 → [H3N:Ag:NH3]+
The result of this reaction is the formation of an ammonia–silver adduct.
In reactions between Lewis acids and bases, there is the formation of an adduct when the highest occupied molecular orbital (HOMO) of a molecule, such as NH3 with available lone electron pair(s) donates lone pairs of electrons to the electron-deficient molecule's lowest unoccupied molecular orbital (LUMO) through a co-ordinate covalent bond; in such a reaction, the HOMO-interacting molecule acts as a base, and the LUMO-interacting molecule acts as an acid. In highly-polar molecules, such as boron trifluoride (BF3), the most electronegative element pulls electrons towards its own orbitals, providing a more positive charge on the less-electronegative element and a difference in its electronic structure due to the axial or equatorial orbiting positions of its electrons, causing repulsive effects from lone pair – bonding pair (Lp–Bp) interactions between bonded atoms in excess of those already provided by bonding pair – bonding pair (Bp–Bp) interactions. Adducts involving metal ions are referred to as co-ordination compounds.
Other acid–base theories 
Lux–Flood definition 
This acid–base theory was a revival of oxygen theory of acids and bases, proposed by German chemist Hermann Lux in 1939, further improved by Håkon Flood circa 1947 and is still used in modern geochemistry and electrochemistry of molten salts. This definition describes an acid as an oxide ion (O2−) acceptor and a base as an oxide ion donor. For example:
- MgO (base) + CO2 (acid) → MgCO3
- CaO (base) + SiO2 (acid) → CaSiO3
3 (base) + S2O2−
7 (acid) → NO+
2 + 2 SO2−
Pearson definition 
In 1963, Ralph Pearson proposed an advanced qualitative concept known as Hard Soft Acid Base principle, later made quantitative with help of Robert Parr in 1984. 'Hard' applies to species that are small, have high charge states, and are weakly polarizable. 'Soft' applies to species that are large, have low charge states and are strongly polarizable. Acids and bases interact, and the most stable interactions are hard–hard and soft–soft. This theory has found use in organic and inorganic chemistry.
Usanovich definition 
Mikhail Usanovich developed a general theory that does not restrict acidity to hydrogen-containing compounds, but his approach, published in 1938, was even more general than Lewis theory. Usanovich's theory can be summarized as defining an acid as anything that accepts negative species or donates positive ones, and a base as the reverse. This defined the concept of redox (oxidation-reduction) as a special case of acid-base reactions
Some examples of Usanovich acid-base reactions include:
- Na2O (base) + SO3 (acid) → 2 Na+ + SO2−
4 (species exchanged: anion O2−)
- 3 (NH4)2S (base) + Sb2S3 (acid) → 6 NH+
4 + 2 SbS3−
4 (species exchanged: anion S2−)
- Na (base) + Cl (acid) → Na+ + Cl− (species exchanged: electron)
Acid–alkali reaction 
An acid–alkali reaction is a special case of an acid–base reaction, where the base used is also an alkali. When an acid reacts with an alkali it forms a metal salt and water. Acid–alkali reactions are also a type of neutralization reaction.
In general, acid–alkali reactions can be simplified to
by omitting spectator ions.
Acids are in general pure substances that contain hydrogen ions (H+) or cause them to be produced in solutions. Hydrochloric acid (HCl) and sulfuric acid (H2SO4) are common examples. In water, these break apart into ions:
- HCl → H+(aq) + Cl−(aq)
- H2SO4 → H+(aq) + HSO−
To produce hydroxide ions in water, the alkali breaks apart into ions as below:
- NaOH → Na+(aq) + OH−(aq)
See also 
- Electron configuration
- Lewis structure
- Resonance structure
- Protonation and Deprotonation
- Nucleophilic substitution and Redox reactions
- Acid–base titration
- Miessler, G.L., Tarr, D. A., "Inorganic Chemistry" (1991) p. 166 – Table of discoveries attributes Antoine Lavoisier as the first to posit a scientific theory in relation to oxyacids.
- Hall, Norris F. (March 1940). "Systems of Acids and Bases". J. Chem. Educ. 17 (3): 124–128. Bibcode:1940JChEd..17..124H. doi:10.1021/ed017p124.
- Miessler, G. L., Tarr, D. A., (1991)
- Meyers, R. (2003). The Basics of Chemistry. Greenwood Press. p. 156. "Inorganic Chemistry" 2nd ed. Pearson Prentice-Hall p. 166 – table of discoveries attributes Justus von Liebig's publication as 1838
- H. L. Finston and A. C. Rychtman, A New View of Current Acid-Base Theories, John Wiley & Sons, New York, 1983, pp. 140–146.
- Miessler, G. L., Tarr, D. A., (1991) "Inorganic Chemistry" 2nd ed. Pearson Prentice-Hall p. 165
- Murray, K. K., Boyd, R. K., et al. (2006) "Standard definition of terms relating to mass spectrometry recommendations" International Union of Pure and Applied Chemistry. – Please note that, in this document, there is no reference to deprecation of "oxonium", which is also still accepted as it remains in the IUPAC Gold book, but rather reveals preference for the term "Hydronium".
- International Union of Pure and Applied Chemistry (2006) IUPAC Compendium of Chemical Terminology, Electronic version Retrieved from International Union of Pure and Applied Chemistry on 9 May 2007 on URL http://goldbook.iupac.org/O04379.html "Oxonium Ions"
- More recent IUPAC recommendations now suggest the newer term "hydronium" be used in favor of the older accepted term "oxonium" to illustrate reaction mechanisms such as those defined in the Brønsted–Lowry and solvent system definitions more clearly, with the Arrhenius definition serving as a simple general outline of acid–base character.
- LeMay, Eugene (2002). Chemistry. Upper Saddle River, New Jersey: Prentice-Hall. p. 602. ISBN 0-13-054383-7.
- Germann, Albert F. O. (6 October 1925). "A General Theory of Solvent Systems". J.Am.Chem.Soc. 47 (10): 2461–2468. doi:10.1021/ja01687a006.
- The term solvonium has replaced the older term lyonium.
- the term solvate has replaced the older term lyate.
- Miessler, G. L., Tarr, D. A., (1991) "Inorganic Chemistry" 2nd ed. Pearson Prentice-Hall pp. 167–169 – According to this page, the original definition was that "acids have a tendency to lose a proton"
- Clayden, J., Warren, S., et al. (2000) "Organic Chemistry" Oxford University Press pp. 182–184
- Miessler, L. M., Tar, D. A., (1991) "Inorganic Chemistry" 2nd ed. Pearson Prentice-Hall p. 166 – Table of discoveries attributes the date of publication/release for the Lewis theory as 1923.
- Miessler, G. L., Tarr, D. A., (1991) "Inorganic Chemistry" 2nd ed. Pearson Prentice-Hall pp. 170–172
- Franz, H. (1966). "Solubility of Water Vapor in Alkali Borate Melts". J. Am. Ceram. Soc. 49 (9): 473–477. doi:10.1111/j.1151-2916.1966.tb13302.x.
- Lux, Hermann (1939). ""Säuren" und "Basen" im Schmelzfluss: die Bestimmung. der Sauerstoffionen-Konzentration". Ztschr. Elektrochem 45 (4): 303–309.
- Flood, H.; Forland, T. (1947). "The Acidic and Basic Properties of Oxides". Acta Chem. Scand. 1 (6): 592–604. doi:10.3891/acta.chem.scand.01-0592. PMID 18907702.
- Drago, Russel S.; Whitten, Kenneth W. (1966). "The Synthesis of Oxyhalides Utilizing Fused-Salt Media". Inorg. Chem. 5 (4): 677–682. doi:10.1021/ic50038a038.
- Pearson, Ralph G. (1963). "Hard and Soft Acids and Bases". J. Am. Chem. Soc. 85 (22): 3533–3539. doi:10.1021/ja00905a001.