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{{Acids_and_Bases}}
An '''acid-base reaction''' is a [[chemical reaction]], that occurs between an [[acid]] and a [[base (chemistry)|base]]. Several concepts that provide alternative definitions for the reaction mechanisms involved and their application in solving related problems exist. 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 [[France|French]] [[chemist]] [[Antoine Lavoisier]], circa 1776.<ref name="lavoisier_1">Miessler, L. M., Tar, D. A., (1991) p166 - Table of discoveries attributes Antoine Lavoisier as the first to posit a scientific theory in relation to [[oxyacid]]s.</ref>

==Historic acid-base theories==
=== Lavoisier's oxygen theory of acids ===
The first scientific concept of acids and bases was provided by Antoine Lavoisier circa 1776. Since Lavoisier's knowledge of [[strong acids]] was mainly restricted to [[oxoacid]]s, such as {{chem|link=nitric acid|HNO|3|}} (nitric acid) and {{chem|link=sulfuric acid|H|2|SO|4}} (sulfuric acid), which tend to contain central atoms in high [[oxidation number|oxidation states]] surrounded by oxygen, and since he was not aware of the true composition of the [[hydrogen halide|hydrohalic acids]] ([[Hydrogen fluoride|HF]], [[HCl]], [[HBr]] (hydrogen fluroide), and [[Hydrogen iodide|HI]]) (hydrogen iodide), he defined acids in terms of their containing ''[[oxygen]]'', which in fact he named from Greek words meaning "acid-former" (from the [[Greek language|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 [[Humphry Davy|Sir Humphry Davy]] in which he proved the lack of oxygen in [[hydrogen sulfide|H<sub>2</sub>S]], [[hydrogen telluride|H<sub>2</sub>Te]], and the [[hydrogen halide|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".<ref name=review1940>{{Cite journal|title=Systems of Acids and Bases|last=Hall|first=Norris F.|journal=J. Chem. Educ.|date=March 1940|volume=17|issue=3|doi=10.1021/ed017p124|pages=124–128}}</ref> 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===
This definition was proposed by [[Justus von Liebig]] circa 1838,<ref name="liebig_1">Miessler, L. M., Tar, D. A., (1991) "''Inorganic Chemistry''" 2<sup>nd</sup> ed. Pearson Prentice-Hall p166 - table of discoveries attributes Justus von Liebig's publication as 1838</ref> based on his extensive works on the chemical composition of [[organic acid]]s. This finished the doctrinal shift from oxygen-based acids to hydrogen-based acids, started by Davy. According to Liebig, an acid is a hydrogen-containing substance in which the hydrogen could be replaced by a metal.<ref name=meyers_156>{{Cite book|author=Meyers, R.|year=2003|title=The Basics of Chemistry|publisher=Greenwood Press|page=156}}</ref> Liebig's definition, while completely empirical, remained in use for almost 50 years until the adoption of the Arrhenius definition.<ref name="liebig_2">H. L. Finston and A. C. Rychtman, A New View of Current Acid-Base Theories, John Wiley & Sons, New York, 1982, pp. 140-146.</ref>

==Common acid-base theories==
=== Arrhenius definition ===
[[File:Arrhenius2.jpg|thumb|right|Svante Arrhenius]]
The Arrhenius definition of acid-base reactions is a development of the hydrogen theory of acids, devised by [[Svante Arrhenius]], which was used to provide a modern definition of acids and bases that followed from his work with [[Friedrich Wilhelm Ostwald]] in establishing the presence of ions in aqueous solution in 1884, and led to Arrhenius receiving the [[Nobel Prize in Chemistry]] in 1903 for "recognition of the extraordinary services... rendered to the advancement of chemistry by his electrolytic theory of dissociation".<ref name="miessler_165">Miessler, L. M., Tar, D. A., (1991) "''Inorganic Chemistry''" 2<sup>nd</sup> ed. Pearson Prentice-Hall p165</ref>

As defined by Arrhenius, acid-base reactions are characterized by Arrhenius acids, which [[dissociation constant|dissociate]] in aqueous solution to form hydrogen ions ({{chem|H|+}}),<ref name="miessler_165" /> and Arrhenius bases, which form hydroxide ({{chem|OH|−}}) ions. More recent [[IUPAC]] recommendations now suggest the newer term "hydronium"<ref>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".</ref> be used in favor of the older accepted term "oxonium"<ref name="iupac_gold">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''"</ref> 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.<ref name="miessler_165" /> The Arrhenius definition can be summarised as "Arrhenius acids form hydrogen ions in aqueous solution with Arrhenius bases forming hydroxide ions."

The ''universal aqueous acid-base definition'' of the Arrhenius concept is described as the formation of water from hydrogen and hydroxide ions, or hydrogen ions and hydroxide ions from the dissociation of an acid and base in aqueous solution:

:{{chem|H|+}} (aq) + {{chem|OH|−}} (aq) {{eqm}} {{chem|H|2|O}}

(In modern times, the use of {{chem|H|+}} is regarded as a shorthand for [[hydronium|{{chem|H|3|O|+}}]], since it is now known that the bare proton {{chem|H|+}} does not exist as a free species in solution.)

This leads to the definition that in Arrhenius acid-base reactions, a salt and water is formed from the reaction between an acid and a base.<ref name="miessler_165" /> In other words, this is a [[Neutralization (chemistry)|neutralization reaction]].

:acid<sup>+</sup> + base<sup>−</sup> → salt + water

The positive ion from a base forms a salt with the negative ion from an acid. For example, two [[mole (unit)|moles]] of the base [[sodium hydroxide]] (NaOH) can combine with one mole of sulfuric acid ({{chem|H|2|SO|4}}) to form two moles of [[water]] and one mole of [[sodium sulfate]].

:2 NaOH + {{chem|H|2|SO|4}} → 2 {{chem|H|2|O}} + {{chem|Na|2|SO|4}}

The Arrhenius definitions of acidity and alkalinity are restricted to [[aqueous solution]]s, and refer to the concentration of the solvent ions. Under this definition, pure {{chem|H|2|SO|4}} or HCl dissolved in toluene are not acidic, and molten KOH and solutions of sodium amide in liquid ammonia are not alkaline.

===Solvent system definition===
One of the limitations of Arrhenius definition was its reliance on water solutions. E. C. Franklin studied the acid-base reactions in liquid ammonia in 1905 and pointed out the similarities to water-based Arrhenius theory, and Albert F. O. Germann, working with liquid {{chem|COCl|2}}, generalized Arrhenius definition to cover aprotic solvents and formulated the solvent system theory in 1925.<ref>{{Cite journal|title=A General Theory of Solvent Systems|journal=J.Am.Chem.Soc.|first=Albert F. O.|last=Germann|date=6 October 1925|volume=47|issue=10|pages=2461–2468|doi=10.1021/ja01687a006}}</ref>

Germann pointed out that in many solvents there is a certain concentration of a positive species, '''solvonium''' (earlier lyonium) cations and negative species, '''solvate''' (earlier lyate) anions, in equilibrium with the neutral solvent molecules. For example, [[water]] and [[ammonia]] undergo such dissociation into [[hydronium]] and [[hydroxide]], and [[ammonium]] and [[amide]], respectively:

: 2 {{chem|H|2|O}} {{eqm}} {{chem|H|3|O|+}} + {{chem|OH|−}}
: 2 {{chem|NH|3|}} {{eqm}} {{chem|NH|4|+}} + {{chem|NH|2|−}}

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]].

: {{chem|N|2|O|4}} {{eqm}} {{chem|NO|+}} + {{chem|NO|3|−}}
: 2 {{chem|SbCl|3}} {{eqm}} {{chem|SbCl|2|+}} + {{chem|SbCl|4|−}}
: {{chem|COCl|2}} {{eqm}} {{chem|COCl|+}} + {{chem|Cl|-}}
A solute causing an increase in the concentration of the solvonium ions and a decrease in the solvate ions is defined as an acid and one causing the reverse is defined as a base. Thus, in liquid [[ammonia]], {{chem|KNH|2}} (supplying {{chem|NH|2|-}}) is a strong base, and {{chem|NH|4|NO|3}} (supplying {{chem|NH|4|+}}) is a strong acid. In liquid [[sulfur dioxide]] ({{chem|SO|2}}), [[thionyl]] compounds (supplying {{chem|SO|2+}}) behave as acids, and [[sulfites]] (supplying {{chem|SO|3|2−}}) behave as bases.

The non-aqueous acid-base reactions in liquid ammonia are similar to the reactions in water:
: 2 {{chem|NaNH|2}} (base) + {{chem|Zn(NH|2|)|2}} ([[amphiphilic]] amide) → {{chem|Na|2|[Zn(NH|2|)|4|]}}
: 2 {{chem|NH|4|I}} (acid) + {{chem|Zn(NH|2|)|2}} (amphiphilic amide) → {{chem|[Zn(NH|3|)|4|)]I|2}}
Nitric acid can be a base in liquid sulfuric acid:
: {{chem|HNO|3}} (base) + 2 {{chem|H|2|SO|4}} → {{chem|NO|2|+}} + {{chem|H|3|O|+}} + 2 {{chem|HSO|4|−}}
The unique strength of this definition shows in describing the reactions in aprotic solvents, for example in liquid {{chem|N|2|O|4}}:
: {{chem|AgNO|3}} (base) + NOCl (acid) → {{chem|N|2|O|4}} (solvent) + AgCl (salt)

Since solvent-system definition depends on the solvent as well as on the compound itself, the same compound can change its role depending on the choice of the solvent. Thus, {{chem|HClO|4}} is a strong acid in water, a weak acid in acetic acid, and a weak base in fluorosulfonic acid. This was seen as both a strength and a weakness, since some substances, such as {{chem|SO|3}} and {{chem|NH|3}} were felt to be acidic or basic on their own right. On the other hand, solvent system theory was criticized as too general to be useful; it was felt that there was something intrinsically acidic about hydrogen compounds, not shared by non-hydrogenic solvonium salts.<ref name=review1940/>

===Brønsted-Lowry definition===
{{Main|Brønsted–Lowry acid-base theory}}

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 [[deprotonation|de-protonation]] of acids—that is, the ability of acids to "donate" hydrogen ions (H<sup>+</sup>) or [[proton (physics)|protons]] to bases, which "accept" them.<ref name="miessler_167">Miessler, L. M., Tar, D. A., (1991) "''Inorganic Chemistry''" 2<sup>nd</sup> ed. Pearson Prentice-Hall p167-169 -- According to this page, the original definition was that "acids have a tendency to lose a proton"</ref> 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.<ref name="miessler_165" /><ref name="miessler_167" /> 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 (chemistry)|neutralization]] is thus absent.<ref name=review1940/>

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.<ref name="Clayden_1">Clayden, J., Warren, S., et al. (2000) "''Organic Chemistry''" Oxford University Press p182-184</ref> 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 ({{chem|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 {{chem|H|+}} from [[hydrochloric acid]] (HCl) produces the chloride ion ({{chem|Cl|−}}), the conjugate base of the acid:

:HCl → {{chem|H|+}} + {{chem|Cl|−}}

The addition of {{chem|H|+}} to the hydroxide ion ({{chem|OH|−}}), a base, produces water ({{chem|H|2|O}}), its conjugate acid:

:{{chem|H|+}} + {{chem|OH|−}} → {{chem|H|2|O}}

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:
:{{chem|H|+}} + {{chem|NH|3}} → {{chem|NH|4|+}}

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:

:{{chem|CH|3|COOH}} + {{chem|NH|3}} → {{chem|NH|4|+}} + {{chem|CH|3|COO|−}}

This definition also explains the dissociation of water into low concentrations of hydronium and hydroxide ions:

:{{chem|H|2|O}} + {{chem|H|2|O}} {{eqm}} {{chem|H|3|O|+}} + {{chem|OH|−}}

Water, being [[amphoteric]], can act as both an acid and a base; here, one molecule of water acts as an acid, donating a {{chem|H|+}} ion and forming the conjugate base, {{chem|OH|−}}, and a second molecule of water acts as a base, accepting the {{chem|H|+}} ion and forming the conjugate acid, {{chem|H|3|O|+}}.

Acid dissociation and acid hydrolysis are seen to be entirely similar phenomena:

: HCl (acid) + {{chem|H|2|O}} (base) {{eqm}} {{chem|H|3|O|+}} (acid) + {{chem|Cl|-}} (base)
: {{chem|NH|4|+}} (acid) + {{chem|H|2|O}} (base) {{eqm}} {{chem|H|3|O|+}} (acid) + {{chem|NH|3}} (base)

as are basic dissociation and basic hydrolysis:

: {{chem|NH|3}} (base) + {{chem|H|2|O}} (acid) {{eqm}} {{chem|NH|4|+}} (acid) + {{chem|OH|-}} (base)
: {{chem|CH|3|COO|-}} (base) + {{chem|H|2|O}} (acid) {{eqm}} {{chem|CH|3|COOH}} (acid) + {{chem|OH|-}} (base)

[[File:Bronsted-lowry-3d-explanation-diagram.png|600px|center]]

Thus, the general formula for acid-base reactions according to the Brønsted-Lowry definition is:

:AH + B → {{chem|BH|+}} + {{chem|A|−}}

where AH represents the acid, B represents the base, {{chem|BH|+}} represents the conjugate acid of B, and {{chem|A|−}} represents the conjugate base of AH.

Although Brønsted-Lowry calls hydrogen-containing substances like HCl acids, KOH and {{chem|KNH|2}} are not bases but salts containing the bases {{chem|OH|-}} and {{chem|NH|2|-}}. Also, some substances, which many chemists considered to be acids, such as {{chem|SO|3}} or {{chem|BCl|3}}, are excluded from this classification due to lack of hydrogen. Gilbert Lewis wrote in 1938, "To restrict the group of acids to those substances which contain hydrogen interferes as seriously with the systematic understanding of chemistry as would the restriction of the term oxidizing agent to substances containing oxygen."<ref name=review1940/>

===Lewis definition===
{{See|Lewis acids and bases}}
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,<ref name="lewis_1">Miessler, L. M., Tar, D. A., (1991) "''Inorganic Chemistry''" 2<sup>nd</sup> ed. Pearson Prentice-Hall p166 - Table of discoveries attributes the date of publication/release for the Lewis theory as 1923.</ref> in the same year as Brønsted-Lowry, but it was not elaborated by him until 1938.<ref name=review1940/> 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.<ref name="lewis_2">Miessler, L. M., Tar, D. A., (1991) "''Inorganic Chemistry''" 2<sup>nd</sup> ed. Pearson Prentice-Hall p170-172</ref>

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) → {{chem|H|2|O}} (l) + NaCl (aq)

The Lewis definition does not regard this reaction as the formation of salt and water or the transfer of {{chem|H|+}} from HCl to {{chem|OH|−}}. Instead, it regards the acid to be the {{chem|H|+}} ion itself, and the base to be the {{chem|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 {{chem|OH|−}} to the {{chem|H|+}} ion. This forms a covalent bond between {{chem|H|+}} and {{chem|OH|−}}, thus producing water ({{chem|H|2|O}}).

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:

: {{chem|Ag|+}} + 2 :{{chem|NH|3}} → {{chem|[H|3|N}}:Ag:{{chem|NH|3|]|+}}

The result of this reaction is the formation of an ammonia-silver adduct.

In reactions between [[Lewis acid]]s and [[Lewis base|bases]], there is the formation of an adduct<ref name="lewis_2" /> when the highest occupied molecular orbital ([[HOMO]]) of a molecule, such as {{chem|NH|3}} 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.<ref name="lewis_2" /> In highly-polar molecules, such as [[boron trifluoride]] ({{chem|BF|3}}),<ref name="lewis_2" /> the most [[electronegativity|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.<ref name="lewis_2" /> Adducts involving metal ions are referred to as co-ordination compounds.<ref name="lewis_2" />

==Other acid-base theories==
=== Usanovich definition ===
Simultaneously with Lewis, a Soviet chemist ''Mikhail Usanovich'' from Tashkent, 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.<ref name=review1940/> 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 pushed the concept of acid-base reactions to its logical limits, and even redefined the concept of [[redox]] (oxidation-reduction) as a special case of acid-base reactions, and so did not become wide spread, despite being easier to understand than Lewis theory, which required detailed familiarity with atomic structure.
Some examples of Usanovich acid-base reactions include:

: {{chem|Na|2|O}} (base) + {{chem|SO|3}} (acid) → 2 {{chem|Na|+}} + {{chem|SO|4|2-}} (species exchanged: anion {{chem|O|2-}})
: 3 {{chem|(NH|4|)|2|S}} (base) + {{chem|Sb|2|S|3}} (acid) → 6 {{chem|NH|4|+}} + 2 {{chem|SbS|4|2-}} (species exchanged: anion {{chem|S|2-}})
: Na (base) + Cl (acid) → {{chem|Na|+}} + {{chem|Cl|-}} (species exchanged: electron)

===Lux-Flood definition===
This acid-base theory was a revival of oxygen theory of acids and bases, proposed by German chemist [[Hermann Lux]]<ref>{{Cite journal|title=Solubility of Water Vapor in Alkali Borate Melts|last=Franz|first=H.|year=1966|journal=J. Am. Ceram. Soc.|volume=49|issue=9|pages=473–477|doi=10.1111/j.1151-2916.1966.tb13302.x}}</ref><ref name=lux>
{{Cite journal|title="Säuren" und "Basen" im Schmelzfluss: die Bestimmung. der Sauerstoffionen-Konzentration|first=Hermann|last=Lux|authorlink=Hermann Lux|journal=Ztschr. Elektrochem|year=1939|volume=45|issue=4|pages=303–309}}</ref> in 1939, further improved by [[Håkon Flood]] circa 1947<ref name=flood>{{Cite journal|title=The Acidic and Basic Properties of Oxides|last=Flood|first=H.|authorlink=Håkon Flood|coauthors=Forland, T.|journal=Acta Chem. Scand.|year=1947|volume=1|pages=592|doi=10.3891/acta.chem.scand.01-0592|pmid=18907702|issue=6}}</ref> and is still used in modern [[geochemistry]] and [[electrochemistry]] of [[molten salt]]s. This definition describes an '''acid''' as an oxide ion ({{chem|O|2-}}) acceptor and a '''base''' as an oxide ion donor. For example:<ref name=drago>{{Cite journal|title=The Synthesis of Oxyhalides Utilizing Fused-Salt Media|first=Russel S.|last=Drago|coauthors=Whitten, Kenneth W.|journal=[[Inorg. Chem.]]|year=1966|volume=5|issue=4|pages=677–682|doi=10.1021/ic50038a038}}</ref>

: MgO (base) + CO<sub>2</sub> (acid) → MgCO<sub>3</sub>
: CaO (base) + SiO<sub>2</sub> (acid) → CaSiO<sub>3</sub>
: {{chem|NO|3|-}} (base) + {{chem|S|2|O|7|2-}} (acid) → {{chem|NO|2|+}} + 2 {{chem|SO|4|2-}}

===Pearson definition===
{{Main|HSAB theory}}
In 1963<ref name=pearson>{{Cite journal|title=Hard and Soft Acids and Bases|last=Pearson|first=Ralph G.|journal= [[J. Am. Chem. Soc.]] |year=1963| volume= 85 |issue=22|pages=3533–3539|doi=10.1021/ja00905a001}}</ref> [[Ralph Pearson]] proposed an advanced qualitative concept known as [[HSAB concept|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.

==Acid-alkali reaction==
{{cleanup|section|date=December 2009}}

{{Merge to|Neutralization (chemistry)|date=January 2010}}
An acid-alkali reaction is a special case of an acid-base reaction, where the [[base (chemistry)|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 (chemistry)|neutralization reaction]].

In general acid-alkali reactions can be simplified to

:[[hydroxide|OH<sup>−</sup>(aq)]] + [[Hydronium|H<sup>+</sup>(aq)]] → [[water (molecule)|H<sub>2</sub>O]]

by omitting [[spectator ions]].

'''Acids''' are generally pure substances which contain hydrogen ions (H<sup>+</sup>) or cause them to be produced in solutions. Hydrochloric acid (HCl) and sulfuric acid (H<sub>2</sub>SO<sub>4</sub>) are common examples. In water, these break apart into ions:

:HCl → H<sup>+</sup>(aq) + Cl<sup>−</sup>(aq)
:H<sub>2</sub>SO<sub>4</sub> → H<sup>+</sup>(aq) + HSO<sub>4</sub><sup>−</sup>(aq)

An '''alkali''' is a base, more precisely a base which contains a metal from column 1 or 2 of the periodic table (the alkali metals or the alkaline earth metals). The Science Module 3 within The Digital Brain [http://www.digitalbrain.com/document.server/subjects/ks3sci/su3/mod3/acidmet.htm] (Science KS3 SU3 Module 3, Acids and Bases) and Doc Brown (http://www.docbrown.info/, 2000-2008, p.&nbsp;3) define alkalis as ''soluble bases'', which means they must be able to dissolve in water. Bases generally are defined as substances which contain hydroxide ion (OH<sup>-</sup>) or produce it in solution. Therefore, we may also speak of hydroxide bases which dissolve in water, and thus these would also be alkalis. Some examples, then, of alkalis would be sodium hydroxide (NaOH), potassium hydroxide (KOH), magnesium hydroxide (Mg(OH)<sub>2</sub>), and calcium hydroxide (Ca(OH)<sub>2</sub>). Note that only hydroxides with an alkali metal—column 1—are very soluble in water; hydroxides with an alkaline earth metal—column 2—are not as soluble. Some sources will even say the alkaline earth metal hydroxides are insoluble.

To produce hydroxide ions in water, the alkali breaks apart into ions as below:

:NaOH → Na<sup>+</sup>(aq) + OH<sup>-</sup>(aq)

However, alkalies may also have a broader definition which includes carbonates (CO<sub>3</sub><sup>2-</sup>) bonded to a column 1 metal, an ammonium ion (NH<sub>4</sub><sup>+</sup>), or an amine (NH<sub>x</sub> radical) as the positive ion. Examples of alkalis would then also include Li<sub>2</sub>CO<sub>3</sub>, Na<sub>2</sub>CO<sub>3</sub>, and (NH<sub>4</sub>)<sub>2</sub>CO<sub>3</sub>.

There seems to be conflicting information on acid-base reactions being '''neutralization reactions'''. Some sources define a neutralization reaction as the reaction between an acid and a base which produces a salt and water. Yet in the book '''Chemical Misconceptions: Prevention, Diagnosis and Cure''' by K. Tabor (2002), it is noted that “the term neutralization is usually reserved for acid-alkali reactions.” Thus this does not make acid-alkali a ''type'' of neutralization reaction, but the ''only kind'' of neutralization reaction.

There are many '''uses''' of neutralization reactions which are acid-alkali reactions. A very common use is antacid tablets. These are designed to neutralize excess stomach acid (HCl) which may be causing discomfort in the stomach or lower esophagus. Also in the digestive tract, neutralization reactions are used when food is moved from the stomach to the intestines. In order for the nutrients to be absorbed through the intestinal wall, an alkaline environment is needed, so the pancreas produce an antacid bicarbonate to cause this transformation to occur. (http://www.wddty.com/UtilityPages/Print.aspx?nodeId=-3363800369331166395)

Another common use, though perhaps not as widely known, is in fertilizers and control of soil pH. Slaked lime (calcium hydroxide) or limestone (calcium carbonate) may be worked into soil that is too acidic for plant growth. (http://www.practicalchemistry.org/experiments/intermediate/acids-alkalis-and-salts/neutralisation-curing-acidity,103,EX.html) Fertilizers which improve plant growth are made by neutralizing sulfuric acid (H<sub>2</sub>SO<sub>4</sub>) or nitric acid (HNO<sub>3</sub>) with ammonia gas (NH<sub>3</sub>) making ammonium sulfate or ammonium nitrate. These are salts utilized in the fertilizer. (http://www.docbrown.info, 2000-2008, p.&nbsp;3)

Industrially, a by-product of the burning of coal, sulfur dioxide gas may combine with water vapor in the air to eventually produce sulfuric acid, which falls as acid rain. To prevent the sulfur dioxide from being released, a device known as a scrubber gleans the gas from smoke stacks. This device first blows calcium carbonate into the combustion chamber where it decomposes into calcium oxide (lime) and carbon dioxide. This lime then reacts with the sulfur dioxide produced forming calcium sulfite. A suspension of lime is then injected into the mixture to produce a slurry, which removes the calcium sulfite and any remaining unreacted sulfur dioxide. (Zumdahl, 2000, p.&nbsp;226, 228)

==See also==
* [[Alkali]]
* [[Electron configuration]]
* [[Lewis structure]]
* [[Resonance structure]]
* [[Protonation]] and [[Deprotonation]]
* [[Nucleophilic substitution]] and [[Redox reaction]]s
* [[Acid-base titration]]

==References==
{{Reflist|colwidth=35em}}

==External links==
* [http://www.anaesthesiamcq.com/AcidBaseBook/ABindex.php Acid-base Physiology: an on-line text]
* [http://users.rcn.com/jkimball.ma.ultranet/BiologyPages/A/Acids_Bases.html John W. Kimball's online Biology book section of acid and bases.]
* [http://dbhs.wvusd.k12.ca.us/webdocs/AcidBase/Early-Acid-Base.html Lavoisier, Davy, and Liebig theories at the ChemTeam Tutorials]
{{Use dmy dates|date=September 2010}}

{{DEFAULTSORT:Acid–Base Reaction}}
[[Category:Acid-base chemistry]]
[[Category:Equilibrium chemistry]]
[[Category:Acids]]
[[Category:Bases]]

[[ar:تفاعل حمض-قلوي]]
[[an:Reacción aceto-base]]
[[de:Säure-Base-Konzepte]]
[[el:Αντίδραση οξέος-βάσης]]
[[es:Reacción ácido-base]]
[[fr:Réaction acido-basique]]
[[is:Sýru-basa hvarf]]
[[it:Reazione acido-base]]
[[hu:Sav-bázis elméletek]]
[[nl:Zuur-basereactie]]
[[ja:酸と塩基]]
[[no:Syre-base-reaksjoner]]
[[pt:Reação ácido-base]]
[[sr:Теорије киселина и база]]
[[sh:Teorije kiselina i baza]]
[[su:Téori réaksi asam-basa]]
[[fi:Happo-emäsreaktio]]
[[tl:Ganting asido-base]]
[[zh:酸碱理论]]

Revision as of 12:52, 13 November 2010

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