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{{otheruses4|acids in chemistry|the drug|Lysergic acid diethylamide}}
{{redirect|Acidity|the novelette|Acidity (Novelette)}}
{{Acids and Bases}}
An '''acid''' (often represented by the generic formula '''HA''' [H⁺A^-]) is traditionally considered any [[chemical compound]] that, when dissolved in [[water]], gives a solution with a [[hydrogen ion]] [[Activity (chemistry)|activity]] greater than in pure water, i.e. a [[pH]] less than 7.0. That approximates the modern definition of [[Johannes Nicolaus Brønsted]] and [[Martin Lowry]], who independently defined an acid as a compound which donates a [[hydrogen ion]] (H⁺) to another compound (called a [[Base (chemistry)|base]]). Common examples include [[acetic acid]] (in [[vinegar]]) and [[sulfuric acid]] (used in [[car battery|car batteries]]). Acid/base systems are different from [[redox]] reactions in that there is no change in [[oxidation state]].

==Definitions==
{{main|acid-base reaction theories}}
The word "acid" comes from the [[Latin]] ''acidus'' meaning "sour," but in [[chemistry]] the term acid has a more specific meaning. There are four common ways to define an acid:
* '''Arrhenius''': According to this definition developed by the [[Sweden|Swedish]] chemist [[Svante Arrhenius]], an acid is a substance that increases the concentration of hydrogen ions (H⁺), which are carried as [[hydronium]] ions (H₃O⁺) when dissolved in [[water]], while bases are substances that increase the concentration of [[hydroxide]] ions (OH^-). This definition limits acids and bases to substances that can dissolve in water. Around 1800, many [[France|French]] chemists, including [[Antoine Lavoisier]], incorrectly believed that all acids contained [[oxygen]]. Indeed the modern German word for oxygen is ''Sauerstoff'' (lit. sour substance), as is the Afrikaans word for oxygen ''suurstof'', with the same meaning. [[England|English]] chemists, including [[Sir Humphry Davy]] at the same time believed all acids contained hydrogen. Arrhenius used this belief to develop this definition of acid.
* '''[[Brønsted-Lowry]]''': According to this definition, an acid is a [[proton]] ([[hydrogen]] nucleus) donor and a base is a proton acceptor. The acid is said to be dissociated after the proton is donated. An acid and the corresponding base are referred to as [[conjugate acid]]-base pairs. Brønsted and Lowry independently formulated this definition, which includes water-insoluble substances not in the Arrhenius definition.
* '''solvent-system definition''': According to this definition, an acid is a substance that, when dissolved in an autodissociating solvent, increases the concentration of the [[solvonium]] cations, such as H₃O⁺ in water, NH₄⁺ in liquid ammonia, NO⁺ in liquid N₂O₄, SbCl₂⁺ in SbCl₃, etc. Base is defined as the substance that increases the concentration of the [[solvate]] anions, respectively OH^-, NH₂^-, NO₃^-, or SbCl₄^-. This definition extends acid-base reactions to nonaqueous systems and even some aprotic systems, where no [[hydrogen]] nuclei are involved in the reactions. This definition is not absolute, a compound acting as acid in one solvent may act as a base in another.
* '''Lewis''': According to this definition developed by [[Gilbert N. Lewis]], an acid is an electron-pair acceptor and a base is an electron-pair donor. (These are frequently referred to as "[[Lewis acid]]s" and "[[Lewis base]]s," and are [[electrophile]]s and [[nucleophile]]s, respectively, in [[organic chemistry]]; Lewis bases are also [[ligands]] in [[coordination compound|coordination chemistry]].) Lewis acids include substances with no transferable [[protons]] (ie H⁺ hydrogen ions), such as [[iron(III) chloride]], and hence the Lewis definition of an acid has wider application than the Brønsted-Lowry definition. In fact, the term '''Lewis acid''' is often used to exclude protic (Brønsted-Lowry) acids. The Lewis definition can also be explained with [[molecular orbital]] theory. In general, an acid can receive an electron pair in its lowest unoccupied orbital ([[LUMO]]) from the highest occupied orbital ([[HOMO]]) of a base. That is, the HOMO from the base and the LUMO from the acid combine to a bonding molecular orbital.

Although not the most general theory, the Brønsted-Lowry definition is the most widely used definition. The strength of an acid may be understood by this definition by the stability of hydronium and the solvated conjugate base upon dissociation. Increasing or decreasing stability of the conjugate base will increase or decrease the acidity of a compound. This concept of acidity is used frequently for [[organic acid]]s such as [[carboxylic acid]]. The molecular orbital description, where the unfilled proton orbital overlaps with a lone pair, is connected to the Lewis definition.

==Properties==
Bronsted-Lowry acids:
*Are generally sour in taste
*Strong or concentrated acids often produce a stinging feeling on [[mucous membrane]]s
*React to indicators as follows: turn blue [[litmus]] and [[methyl orange]] red, do not change the color of [[phenolphthalein]]
*Will react with metals to produce a metal salt and hydrogen
*Will react with metal carbonates to produce water, CO₂ and a salt
*Will react with a base to produce a salt and water
*Will react with a metal oxide to produce water and a salt
*Will conduct electricity, depending on the degree of dissociation
*Will produce solvonium ions, such as hydronium (H₃O⁺) ions in water
*Will denature proteins

[[Strong acid]]s and many concentrated acids are dangerous, causing severe burns for even minor contact. Acids are corrosive. Generally, acid burns are treated by rinsing the affected area abundantly with running water (15 minutes) and followed up with immediate medical attention. In the case of highly concentrated acids, the acid should first be wiped off as much as possible, otherwise the exothermic mixing of the acid and the water could cause severe thermal burns. Acids may also be dangerous for reasons not related to their acidity, see an appropriate [[MSDS]] for more detailed information.

==Nomenclature==
In the classical naming system, acids are named according to their [[anion]]s. That ionic suffix is dropped and replaced with a new suffix (and sometimes prefix), according to the table below. For example, HCl has [[chloride]] as its anion, so the -ide suffix makes it take the form [[hydrochloric acid]]. In the [[IUPAC]] naming system, "aqueous" is simply added to the name of the ionic compound. Thus, for hydrogen chloride, the IUPAC name would be aqueous hydrogen chloride.

Classical naming system:
{| border="1" cellpadding="4" align="center" cellspacing="0" style="background: #f9f9f9; color: black; border: 1px #aaa solid; border-collapse: collapse;"
!Anion Prefix
!Anion Suffix
!Acid Prefix
!Acid Suffix
!Example
|-
|per
|ate
|per
|ic acid
|[[perchloric acid]] (HClO₄)
|-
|
|ate
|
|ic acid
|[[chloric acid]] (HClO₃)
|-
|
|ite
|
|ous acid
|[[chlorous acid]] (HClO₂)
|-
|hypo
|ite
|hypo
|ous acid
|[[hypochlorous acid]] (HClO)
|-
|
|ide
|hydro
|ic acid
|[[hydrochloric acid]] (HCl)
|}

==Chemical characteristics==
In water the following [[chemical equilibrium|equilibrium]] occurs between a weak acid (HA) and water, which acts as a base:

HA([[Aqueous solution|aq]]) + H₂O {{unicode|⇌}} H₃O⁺(aq) + A^-(aq)

The [[acidity constant]] (or acid dissociation constant) is the equilibrium constant for the reaction of HA with water:

:<math>K_a = {[\mbox{H}_3\mbox{O}^+]\cdot[\mbox{A}^-] \over [\mbox{HA}]}</math>

[[Strong acids]] have large ''K''_a values (i.e. the reaction equilibrium lies far to the right; the acid is almost completely dissociated to H₃O⁺ and A^-). Strong acids include the heavier [[hydrohalic acid]]s: [[hydrochloric acid]] (HCl), [[hydrobromic acid]] (HBr), and [[hydroiodic acid]] (HI). (However, [[hydrofluoric acid]], HF, is relatively weak.) For example, the ''K''_a value for hydrochloric acid (HCl) is 10⁷.

[[Weak acid]]s have small ''K''_a values (i.e. at equilibrium significant amounts of HA and A⁻ exist together in solution; modest levels of H₃O⁺ are present; the acid is only partially dissociated). For example, the K_a value for acetic acid is 1.8 x 10^-5. Most organic acids are weak acids. [[Oxoacid]]s, which tend to contain central atoms in high oxidation states surrounded by oxygen may be quite strong or weak. [[Nitric acid]], sulfuric acid, and [[perchloric acid]] are all strong acids, whereas [[nitrous acid]], [[sulfurous acid]] and [[hypochlorous acid]] are all weak.

Note on terms used:
* The terms "[[hydrogen]] ion" and "proton" are used interchangeably; both refer to H⁺.
* In aqueous solution, the water is protonated to form [[hydronium]] ion, H₃O⁺(aq). This is often abbreviated as H⁺(aq) even though the symbol is not chemically correct.
* The strength of an acid is measured by its [[acid dissociation constant]] (''K''_a) or equivalently its p''K''_a (p''K''_a= - log(''K''_a)).
* The [[pH]] of a solution is a measurement of the concentration of hydronium. This will depend on the concentration and nature of acids and bases in solution.

===Polyprotic acids===
Polyprotic acids are able to donate more than one proton per acid molecule, in contrast to monoprotic acids that only donate one proton per molecule. Specific types of polyprotic acids have more specific names, such as '''diprotic acid''' (two potential protons to donate) and '''triprotic acid''' (three potential protons to donate).

A monoprotic acid can undergo one [[dissociation (chemistry)|dissociation]] (sometimes called ionization) as follows and simply has one acid dissociation constant as shown below:

:::::HA(aq) + H₂O(l) {{unicode|⇌}} H₃O⁺(aq) + A⁻(aq) &nbsp; &nbsp; &nbsp; &nbsp; ''K''_a

A diprotic acid (here symbolized by H₂A) can undergo one or two dissociations depending on the pH. Each dissociation has its own dissociation constant, K_a1 and K_a2.

:::::H₂A(aq) + H₂O(l) {{unicode|⇌}} H₃O⁺(aq) + HA⁻(aq) &nbsp; &nbsp; &nbsp; ''K''_a1

:::::HA⁻(aq) + H₂O(l) {{unicode|⇌}} H₃O⁺(aq) + A²⁻(aq)&nbsp; &nbsp; &nbsp;&nbsp; ''K''_a2

The first dissociation constant is typically greater than the second; i.e., ''K''_a1 > ''K''_a2 . For example, [[sulfuric acid]] (H₂SO₄) can donate one proton to form the [[bisulfate]] anion (HSO₄⁻), for which ''K''_a1 is very large; then it can donate a second proton to form the [[sulfate]] anion (SO₄²⁻), wherein the ''K''_a2 is intermediate strength. The large ''K''_a1 for the first dissociation makes sulfuric a strong acid. In a similar manner, the weak unstable [[carbonic acid]] (H₂CO₃) can lose one proton to form [[bicarbonate]] anion (HCO₃⁻) and lose a second to form [[carbonate]] anion (CO₃²⁻). Both ''K''_a values are small, but ''K''_a1 > ''K''_a2 .

A triprotic acid (H₃A) can undergo one, two, or three dissociations and has three dissociation constants, where ''K''_a1 > ''K''_a2 > ''K''_a3 .

:::::H₃A(aq) + H₂O(l) {{unicode|⇌}} H₃O⁺(aq) + H₂A⁻(aq) &nbsp; &nbsp; &nbsp;&nbsp; ''K''_a1

:::::H₂A⁻(aq) + H₂O(l) {{unicode|⇌}} H₃O⁺(aq) + HA²⁻(aq) &nbsp; &nbsp; &nbsp; ''K''_a2

:::::HA²⁻(aq) + H₂O(l) {{unicode|⇌}} H₃O⁺(aq) + A³⁻(aq) &nbsp; &nbsp;&nbsp;&nbsp;&nbsp;&nbsp; ''K''_a3

An [[inorganic]] example of a triprotic acid is orthophosphoric acid (H₃PO₄), usually just called [[phosphoric acid]]. All three protons can be successively lost to yield H₂PO₄⁻, then HPO₄²⁻, and finally PO₄³⁻ , the orthophosphate ion, usually just called [[phosphate]]. An [[organic compound|organic]] example of a triprotic acid is [[citric acid]], which can successively lose three protons to finally form the [[citrate]] ion. Even though the positions of the protons on the original molecule may be equivalent, the successive ''K''_a values will differ since it is energetically less favorable to lose a proton if the conjugate base is more negatively charged.

===Neutralization===
[[Neutralization]] is the reaction between an acid and a base, producing a [[salt (chemistry)|salt]] and [[water (molecule)|water]]; for example, hydrochloric acid and sodium hydroxide form sodium chloride and water:

::HCl(aq) + NaOH(aq) → H₂O(l) + NaCl(aq)

Neutralization is the basis of [[titration]], where a [[pH indicator]] shows equivalence point when the equivalent number of moles of a base have been added to an acid. It is often wrongly assumed that neutralization should result in a solution with pH 7.0, which is only the case with similar acid and base strengths during a reaction.

Neutralization with an base weaker than the acid results in an weakly acidic salt. An example is the weakly acidic [[ammonium chloride]], which is produced from the strong acid [[hydrogen chloride]] and the weak base [[ammonia]]. Conversely, neutralizing a weak acid with a strong base gives a weakly basic salt, e.g. [[sodium fluoride]] from [[hydrogen fluoride]] and [[sodium hydroxide]].

===Weak acid/weak base equilibria===
{{main|Henderson-Hasselbalch equation}}
In order to lose a proton, it is necessary that the pH of the system rise above the p''K''_a of the protonated acid. The decreased concentration of H⁺ in that basic solution shifts the equilibrium towards the conjugate base form (the deprotonated form of the acid). In lower-pH (more acidic) solutions, there is a high enough H⁺ concentration in the solution to cause the acid to remain in its protonated form, or to protonate its conjugate base (the deprotonated form).

Solutions of weak acids and salts of their conjugate bases form [[buffer solution]]s.

== Applications of acids ==

There are numerous uses for acids. Acids are often used to remove rust and other corrosion from metals in a process known as [[pickling (metal)|pickling]]. They may be used as an electrolyte in a [[wet cell battery]], such as [[sulfuric acid]] in a [[car battery]]. In humans and many other animals, [[hydrochloric acid]] is a part of the [[gastric acid]] secreted within the [[stomach]] to help hydrolyze [[protein]]s and [[polysaccharide]]s, as well as converting the inactive pro-enzyme, [[pepsinogen]] into the enzyme, [[pepsin]]. Acids are used as [[catalyst]]s; for example, [[sulfuric acid]] is used in very large quantities in the [[alkylation]] process to produce gasoline.

== Common acids ==
*[[Citric Acid]]
=== Mineral acids ===
*Solutions of hydrogen halides, such as [[hydrochloric acid]] (HCl) and [[hydrobromic acid]] (HBr)
*[[Sulfuric acid]] (H₂SO₄)
*[[Nitric acid]] (HNO₃)
*[[Phosphoric acid]] (H₃PO₄)
*[[Chromic acid]] (H₂CrO₄)

=== Sulfonic acids ===
*Methanesulfonic acid (aka mesylic acid) (MeSO₃H)
*Ethanesulfonic acid (aka esylic acid) (EtSO₃H)
*Benzenesulfonic acid (aka besylic acid) (PhSO₃H)
*Toluenesulfonic acid (aka tosylic acid, or (C₆H₄(CH₃)(SO₃H))

=== Carboxylic acids ===
*[[Formic acid]]
*[[Acetic acid]]

==References==
{{reflist}}
* [http://www.csudh.edu/oliver/chemdata/data-ka.htm Listing of strengths of common acids and bases]
* Zumdahl, Chemistry, 4th Edition.

== See also ==

; Chemistry
* [[Acid value]]
* [[Acid salt]]
* [[Base (chemistry)]]
* [[Basic salt]]
* [[Binary acid]]
* [[Vitriol]]
* [[Acid-base extraction]]
; Environment

* [[Acid rain]]
* [[Ocean acidification]]

==External links==
* [http://scienceaid.co.uk/chemistry/physical/acidbases.html Science Aid: Acids and Bases] Information for High School students
* [http://www2.iq.usp.br/docente/gutz/Curtipot_.html Curtipot]: Acid-Base equilibria diagrams, [[pH]] calculation and [[titration]] curves simulation and analysis - [[freeware]]
* [http://canadaconnects.ca/chemistry/10081/ A summary of the Properties of Acids for the beginning chemistry student]
*[http://www.unece.org/env/lrtap/ The UN ECE Convention on Long-Range Transboundary Air Pollution]

[[Category:Chemical substances]]
[[Category:Acids| ]]
[[Category:Acid-base chemistry]]

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