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{{Short description|Chemical reaction between an acid and a base}}
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{{Acids and bases}}
In [[chemistry]], an '''acid–base reaction''' is a [[chemical reaction]] that occurs between an [[acid]] and a [[base (chemistry)|base]]. It can be used to determine [[pH]] via [[Acid–base titration|titration]]. Several [[theory|theoretical]] frameworks provide alternative conceptions of the reaction mechanisms and their application in solving related problems; these are called the acid–base theories, for example, [[Brønsted–Lowry acid–base theory]].
Their importance becomes apparent in analyzing acid–base reactions for gaseous or liquid species, or when acid or base character may be somewhat less apparent. The first of these concepts was provided by the French [[chemist]] [[Antoine Lavoisier]], around 1776.<ref name="lavoisier_1">{{harvnb|Miessler|Tarr|1991|p=166}} – Table of discoveries attributes Antoine Lavoisier as the first to posit a scientific theory in relation to [[oxyacid]]s.</ref>
It is important to think of the acid–base reaction models as theories that complement each other.<ref>{{cite journal|last=Paik|first=Seoung-Hey|title=Understanding the Relationship Among Arrhenius, Brønsted–Lowry, and Lewis Theories|journal=Journal of Chemical Education|language=en|volume=92|issue=9|pages=1484–1489|doi=10.1021/ed500891w|bibcode=2015JChEd..92.1484P|year=2015}}</ref> For example, the current Lewis model has the broadest definition of what an acid and base are, with the Brønsted–Lowry theory being a subset of what acids and bases are, and the Arrhenius theory being the most restrictive.
==Acid–base definitions==
===Historic development===
The concept of an acid–base reaction was first proposed in 1754 by [[Guillaume-François Rouelle]], who introduced the word "[[base (chemistry)|base]]" into chemistry to mean a substance which reacts with an acid to give it solid form (as a salt). Bases are mostly bitter in nature.<ref>{{cite journal | author = Jensen, William B.|author1-link=William B. Jensen| title = The origin of the term "base"| journal = The Journal of Chemical Education | year = 2006 | volume = 83 | pages = 1130 | doi = 10.1021/ed083p1130 | issue = 8 | bibcode = 2006JChEd..83.1130J }}</ref>
====Lavoisier's oxygen theory of acids====
The first scientific concept of acids and bases was provided by [[Lavoisier]] in around 1776. Since Lavoisier's knowledge of [[strong acid]]s was mainly restricted to [[oxoacid]]s, such as {{
====Liebig's hydrogen theory of acids====
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As defined by Arrhenius:
* An ''
This causes the [[protonation]] of water, or the creation of the [[hydronium]] (
* An ''
The Arrhenius definitions of [[acidity]] and [[alkalinity]] are restricted to aqueous solutions and are not valid for most non-aqueous solutions, and refer to the concentration of the solvent ions. Under this definition, pure
The reaction of an acid with a base is called a [[neutralization (chemistry)|neutralization]] reaction. The products of this reaction are a [[salt (chemistry)|salt]] and water.
In this traditional representation an acid–base neutralization reaction is formulated as a [[salt metathesis reaction|double-replacement reaction]]. For example, the reaction of [[hydrochloric acid]]
The modifier ([[aqueous solution|aq]]) in this equation was implied by Arrhenius, rather than included explicitly. It indicates that the substances are dissolved in water. Though all three substances, HCl, NaOH and NaCl are capable of existing as pure compounds, in [[aqueous solution]]s they are fully dissociated into the aquated ions {{chem2|H
====Example:
Whether commercially or domestically prepared, the principles behind baking powder formulations remain the same. The acid–base reaction can be generically represented as shown:<ref>{{cite book | editor=A.J. Bent | title=The Technology of Cake Making | edition=6 | year=1997 | page=102 | publisher=Springer | isbn=9780751403497 | url=https://books.google.com/books?id=OTy8aIWxHhQC&pg=PA102 | access-date=2009-08-12}}</ref>
<math chem display=block>\ce{NaHCO3 + H+ -> Na+ + CO2 + H2O}</math>
The real reactions are more complicated because the acids are complicated. For example, starting with sodium bicarbonate and [[monocalcium phosphate]] ({{chem2|Ca(H2PO4)2}}), the reaction produces carbon dioxide by the following [[stoichiometry]]:<ref name="KO">John Brodie, John Godber "Bakery Processes, Chemical Leavening Agents" in Kirk-Othmer Encyclopedia of Chemical Technology 2001, John Wiley & Sons. {{doi|10.1002/0471238961.0308051303082114.a01.pub2}}</ref>
<math chem display=block>\ce{14 NaHCO3 + 5 Ca(H2PO4)2 -> 14 CO2 + Ca5(PO4)3OH + 7 Na2HPO4 + 13 H2O}</math>
[[Image:Calcium dihydrogen phosphate.png|thumb|220px|[[Monocalcium phosphate]] ("MCP") is a common acid component in domestic baking powders.]]
A typical formulation (by weight) could call for 30% sodium bicarbonate, 5–12% [[monocalcium phosphate]], and 21–26% [[sodium aluminium sulfate]]. Alternately, a commercial baking powder might use [[sodium acid pyrophosphate]] as one of the two acidic components instead of sodium aluminium sulfate. Another typical acid in such formulations is [[cream of tartar]] ({{chem2|KC4H5O6}}), a derivative of [[tartaric acid]].<ref name="KO"/>
===Brønsted–Lowry definition===
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|data2= [[Johannes Nicolaus Brønsted]] and [[Martin Lowry|Thomas Martin Lowry]]}}
The Brønsted–Lowry definition, formulated in 1923, independently by [[Johannes Nicolaus Brønsted]] in Denmark and [[Martin Lowry]] in England,<ref>{{cite journal |last= Brönsted |first= J.N. |date= 1923 |title= Einige Bemerkungen über den Begriff der Säuren und Basen |trans-title=Some observations about the concept of acids and bases |journal= Recueil des Travaux Chimiques des Pays-Bas |volume= 42 |issue= 8 |pages= 718–728|doi= 10.1002/recl.19230420815 }}</ref><ref>{{cite journal |last= Lowry |first= T.M. |date= 1923 |url= https://archive.org/stream/ost-chemistry-chemistryindustr01soci/chemistryindustr01soci#page/n65/mode/2up |title= The uniqueness of hydrogen |journal= [[Journal of the Society of Chemical Industry]] |volume= 42 |issue= 3 |pages= 43–47|doi= 10.1002/jctb.5000420302 }}</ref> is based upon the idea of [[protonation]] of bases through the [[deprotonation]] of acids – that is, the ability of acids to "donate" hydrogen ions ({{chem2|H
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">{{harvnb|Clayden|Greeves|Warren|Wothers|
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 [[Peter Debye|Debye]], [[Lars Onsager|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.<ref name=review1940/> Brønsted–Lowry acid–base behavior is formally independent of any solvent, making it more all-encompassing than the Arrhenius model. The calculation of [[pH]] under the Arrhenius model depended on alkalis (bases) dissolving in water ([[aqueous solution]]). The Brønsted–Lowry model expanded what could be [[pH]] tested using insoluble and soluble solutions (gas, liquid, solid).
The general formula for acid–base reactions according to the Brønsted–Lowry definition is:
where HA represents the acid, B represents the base, {{chem2|BH
For example, a Brønsted–Lowry model for the dissociation of [[hydrochloric acid]] (HCl) in [[aqueous solution]] would be the following:
<math chem display=block>\underset{\text{acid}}{\ce{HCl_{\,}}} \ + \ \underset{\text{base}}{\ce{H2O}} \quad \ce{<=>} \quad \underset{\text{conjugate } \atop \text{acid }}{\ce{H3O+}} \ + \underset{\text{conjugate} \atop \text{base}}{\ce{Cl_{\,}-}}</math>
The removal of {{chem2|H
Water is [[amphoterism|amphoteric]]
<math chem display=block>\ce{H2O + H2O <=> H3O+ + OH-}</math>
This equation is demonstrated in the image below:
[[File:Bronsted lowry 3d diagram.png|600px|center]]
Here, one molecule of water acts as an acid, donating an {{chem2|H
As an example of water acting as an acid, consider an aqueous solution of [[pyridine]],
<math chem display=block>\ce{C5H5N + H2O <=> [C5H5NH]+ + OH-}</math>
In this example, a water molecule is split into a hydrogen ion, which is donated to a pyridine molecule, and a hydroxide ion.
In the Brønsted–Lowry model, the solvent does not necessarily have to be water, as is required by the [[#Arrhenius definition|Arrhenius
<math chem display=block>\ce{CH3COOH + NH3 <=> NH4+ + CH3COO-}</math>
An {{chem2|H
The Brønsted–Lowry model calls hydrogen-containing substances (like {{chem2|HCl}}) acids. Thus, some substances, which many chemists considered to be acids, such as
===Lewis definition===
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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">{{harvnb|Miessler|Tarr|1991|p=166}} – Table of discoveries attributes the date of publication/release for the Lewis theory as 1924.</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">{{harvnb|Miessler|Tarr|1991|pp=170–172}}</ref>
For example, [[boron trifluoride]],
<math chem display=block>\ce{BF3 + F- -> BF4-}</math>
is a typical Lewis acid, Lewis base reaction. All compounds of [[boron group|group 13]] elements with a formula
Adducts involving metal ions are referred to as co-ordination compounds; each ligand donates a pair of electrons to the metal ion.<ref name="lewis_2"/> The reaction
<math chem display=block>\ce{[Ag(H2O)4]+ + 2 NH3 -> [Ag(NH3)2]+ + 4 H2O}</math>
can be seen as an acid–base reaction in which a stronger base (ammonia) replaces a weaker one (water).
The Lewis and Brønsted–Lowry definitions are consistent with each other since the reaction
<math chem display=block>\ce{H+ + OH- <=> H2O}</math>
is an acid–base reaction in both theories.
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Germann pointed out that in many solutions, there are ions in equilibrium with the neutral solvent molecules:
* [[onium ion|solvonium ions]]: a generic name for positive ions.
* [[ate complex|solvate ions]]: a generic name for negative ions.
For example, water and [[ammonia]] undergo such dissociation into [[hydronium]] and [[hydroxide]], and [[ammonium]] and [[metal amides#Alkali metal amides|amide]], respectively:
<math chem display=block>\begin{align}
\ce{2 H2O} & \ce{\, <=> H3O+ + OH-} \\[4pt]
\ce{2 NH3} & \ce{\, <=> NH4+ + NH2-}
\end{align}</math>
Some aprotic systems also undergo such dissociation, such as [[dinitrogen tetroxide]] into [[nitrosonium]] and [[nitrate]],{{#tag:ref|Pure N<sub>2</sub>O<sub>4</sub> does not undergo such dissolution. However, it becomes electrically conductive when mixed with a polarized compound, which is believed to correspond with the establishment of such an equilibrium.<ref>{{Greenwood&Earnshaw1st|page=525}}</ref>|group=note}} [[antimony trichloride]] into dichloroantimonium and tetrachloroantimonate, and phosgene into chlorocarboxonium and [[chloride]]:
<math chem display=block>\begin{align}
\ce{N2O4} & \ce{\, <=> NO+ + NO3-} \\[4pt]
\ce{2 SbCl3} & \ce{\, <=> SbCl2+ + SbCl4-} \\[4pt]
\ce{COCl2} & \ce{\, <=> COCl+ + Cl-}
\end{align}</math>
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, {{
The non-aqueous acid–base reactions in liquid ammonia are similar to the reactions in water:
<math chem display=block>\begin{align}
\underset{\text{base}}{\ce{2 NaNH2}} + \underset{\text{amphiphilic} \atop \text{amide}}{\ce{Zn(NH2)2}} &\longrightarrow \ce{Na2[Zn(NH2)4]} \\[4pt]
\underset{\text{acid}}{\ce{2 NH4I}} \ + \ \ce{Zn(NH2)2} &\longrightarrow \ce{[Zn(NH3)4]I2}
\end{align}</math>
Nitric acid can be a base in liquid sulfuric acid:
<math chem display=block>\underset{\text{base}}{\ce{HNO3}} + \ce{2 H2SO4 -> NO2+ + H3O+ + 2 HSO4-}</math>
The unique strength of this definition shows in describing the reactions in aprotic solvents; for example, in liquid {{
<math chem display=block>\underset{\text{base}}{\ce{AgNO3}} + \underset{\text{acid}}{\ce{NOCl_{\ }}} \longrightarrow \underset{\text{solvent}}{\ce{N2O4}} + \underset{\text{salt}}{\ce{AgCl_{\ }}}</math>
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: {{chem2|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 {{chem2|SO3}} and {{chem2|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.<ref name=review1940/>
===Lux–Flood definition===
This acid–base theory was a revival of the oxygen theory of acids and bases proposed by German chemist [[Hermann Lux]]<ref>{{cite journal |last=Franz |first=H. |year=1966 |title=Solubility of Water Vapor in Alkali Borate Melts |journal=[[Journal of the American Ceramic Society]] |volume=49 |issue=9 |pages=473–477 |doi=10.1111/j.1151-2916.1966.tb13302.x}}</ref><ref name="lux">{{cite journal |last=Lux |first=Hermann |author-link=Hermann Lux |year=1939 |title="Säuren" und "Basen" im Schmelzfluss: die Bestimmung. der Sauerstoffionen-Konzentration |journal=[[Zeitschrift für Elektrochemie|Z. Elektrochem
<math chem display=block>\begin{array}{ccccl}
_\text{(base)} & & _\text{(acid)} \\[4pt]
\ce{MgO} &+& \ce{CO2} &\longrightarrow& \ce{MgCO3} \\[4pt]
\ce{CaO} &+& \ce{SiO2} &\longrightarrow& \ce{CaSiO3} \\[4pt]
\ce{NO3-} &+& \ce{S2O7^2-} \!\! &\longrightarrow& \ce{NO2+ + 2 SO4^2-}
\end{array}</math>
This theory is also useful in the systematisation of the reactions of [[noble gas compound]]s, especially the xenon oxides, fluorides, and oxofluorides.<ref>{{Greenwood&Earnshaw1st|page=1056}}</ref>
===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.<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 defined the concept of [[redox]] (oxidation-reduction) as a special case of acid–base reactions.
Some examples of Usanovich acid–base reactions include:
<math chem display=block>\begin{array}{ccccll}
_\text{(base)} & & _\text{(acid)} \\[4pt]
\ce{Na2O} &+& \ce{SO3} &\longrightarrow& \ce{2Na+ {} + \ SO4^2-} & \text{(species exchanged: } \ce{O^2-} \text{anion)} \\[4pt]
\ce{3(NH4)2S} &+& \ce{Sb2S5} &\longrightarrow& \ce{6 NH4+ {} + \ 2 SbS4^3-} & \text{(species exchanged: } \ce{3 S^2-} \text{ anions)} \\[4pt]
\ce{2Na} &+& \ce{Cl2} &\longrightarrow& \ce{2 Na+ {} + \ 2 Cl-} & \text{(species exchanged: 2 electrons)}
\end{array}</math>
==Rationalizing the strength of Lewis acid–base interactions==
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{{Main|ECW model}}
The [[ECW model]] created by [[Russell S. Drago]] is a quantitative model that describes and predicts the strength of Lewis acid base interactions, {{math|−Δ''H''}}. The model assigned
<math display=block>-\Delta H = E_{\rm A}E_{\rm B} + C_{\rm A}C_{\rm B} + W</math>
The
==Acid–base equilibrium==
{{Main|Acid dissociation constant}}
The reaction of a strong acid with a strong base is essentially a quantitative reaction. For example,
In this reaction both the sodium and chloride ions are spectators as the neutralization reaction,
<math chem display=block>\ce{H + OH- -> H2O}</math>
does not involve them. With weak bases addition of acid is not quantitative because a solution of a weak base is a [[buffer solution]]. A solution of a weak acid is also a buffer solution. When a weak acid reacts with a weak base an equilibrium mixture is produced. For example, [[adenine]], written as AH, can react with a hydrogen [[phosphate]] ion, {{
<math chem display=block>\ce{AH + HPO4^2- <=> A- + H2PO4-}</math>
The equilibrium constant for this reaction can be derived from the acid dissociation constants of adenine and of the dihydrogen phosphate ion.
<math chem display=block>\begin{align}
\left[\ce{A-}\right] \! \left[\ce{H+}\right] &= K_{a1}\bigl[\ce{AH}\bigr] \\[4pt]
\left[\ce{HPO4^2-}\right] \! \left[\ce{H+}\right] &= K_{a2}\left[\ce{H2PO4-}\right]
\end{align}</math>
The notation [X] signifies "concentration of X". When these two equations are combined by eliminating the hydrogen ion concentration, an expression for the equilibrium constant,
<math chem display=block>\left[\ce{A-}\right] \! \left[\ce{H2PO4-}\right] = K \bigl[\ce{AH}\bigr] \! \left[\ce{HPO4^2-}\right]; \quad K = \frac{K_{a1}}{K_{a2}}</math>
==Acid–alkali reaction==
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In general, acid–alkali reactions can be simplified to
:<math chem>\ce{OH_{(aq)}- + H+_{(aq)} -> H2O} </math>
by omitting [[spectator ion]]s.
Acids are in general pure substances that contain [[hydron (chemistry)|hydrogen cations]] ({{
:<math chem>\begin{align}
\ce{HCl} &\longrightarrow \ce{H_{(aq)}+ {} + Cl_{(aq)}- } \\[4pt]
\ce{H2SO4} &\longrightarrow \ce{H_{(aq)}+ {} + HSO4_{\,(aq)}- }
\end{align}</math>
The alkali breaks apart in water, yielding dissolved hydroxide ions:
:<math chem>\ce{
==See also==
* [[Acid–base titration]]
* [[Deprotonation]]
* [[Donor number]]
* [[Electron configuration]]
* [[Gutmann–Beckett method]]
* [[Lewis structure]]
* [[Nucleophilic substitution]]
* [[Neutralization (chemistry)]]
* [[Protonation]]
* [[Redox]] reactions
* [[Resonance (chemistry)]]
==Notes==
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===Sources===
*{{cite book |last1=
*{{cite book |first1= H.L. |last1= Finston |first2= A.C. |last2= Rychtman |title= A New View of Current
*{{cite book |last= Meyers |first= R. |year= 2003 |title= The Basics of Chemistry |publisher= Greenwood Press}}
*{{cite book |last1= Miessler |first1= G.L. |last2= Tarr |first2= D.A. |title= Inorganic Chemistry |date= 1991}}
==External links==
*[http://www.anaesthesiamcq.com/AcidBaseBook/ABindex.php
*[https://web.archive.org/web/20070207082349/http://users.rcn.com/jkimball.ma.ultranet/BiologyPages/A/Acids_Bases.html John W. Kimball's online
{{Authority control}}
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[[Category:Acids]]
[[Category:Bases (chemistry)]]
[[Category:
[[Category:Equilibrium chemistry]]
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