Acid

This article is about acids in chemistry. For other uses, seeAcid (disambiguation).

"Acidity" and "acidic" redirect here. For the novelette, seeAcidity(novelette).For the band, seeAcidic (band).

Anacidis amoleculeorioncapable of either donating aproton(i.e. hydrogen ion, H⁺), known as aBrønsted–Lowry acid,or forming acovalent bondwith anelectron pair,known as aLewis acid.

Zinc,a typical metal, reacting withhydrochloric acid,a typical acid

The first category of acids are the proton donors, orBrønsted–Lowry acids.In the special case ofaqueous solutions,proton donors form thehydronium ionH₃O⁺and are known asArrhenius acids.BrønstedandLowrygeneralized the Arrhenius theory to include non-aqueoussolvents.A Brønsted or Arrhenius acid usually contains a hydrogen atom bonded to a chemical structure that is still energetically favorable after loss of H⁺.

Aqueous Arrhenius acids have characteristic properties that provide a practical description of an acid.Acids form aqueous solutions with a sour taste, can turn bluelitmusred, and react withbasesand certain metals (likecalcium) to formsalts.The wordacidis derived from theLatinacidus,meaning 'sour'.An aqueous solution of an acid has apHless than 7 and is colloquially also referred to as "acid" (as in "dissolved in acid" ), while the strict definition refers only to thesolute.A lower pH means a higheracidity,and thus a higher concentration ofpositive hydrogen ionsin the solution. Chemicals or substances having the property of an acid are said to beacidic.

Common aqueous acids includehydrochloric acid(a solution ofhydrogen chloridethat is found ingastric acidin the stomach and activatesdigestive enzymes),acetic acid(vinegar is a dilute aqueous solution of this liquid),sulfuric acid(used incar batteries), andcitric acid(found in citrus fruits). As these examples show, acids (in the colloquial sense) can be solutions or pure substances, and can be derived from acids (in the strictsense) that are solids, liquids, or gases.Strong acidsand some concentrated weak acids arecorrosive,but there are exceptions such ascarboranesandboric acid.

The second category of acids areLewis acids,which form a covalent bond with an electron pair. An example isboron trifluoride(BF₃), whose boron atom has a vacantorbitalthat can form a covalent bond by sharing a lone pair of electrons on an atom in a base, for example the nitrogen atom inammonia(NH₃).Lewisconsidered this as a generalization of the Brønsted definition, so that an acid is a chemical species that accepts electron pairs either directlyorby releasing protons (H⁺) into the solution, which then accept electron pairs. Hydrogen chloride, acetic acid, and most other Brønsted–Lowry acids cannot form a covalent bond with an electron pair, however, and are therefore not Lewis acids.Conversely, many Lewis acids are not Arrhenius or Brønsted–Lowry acids. In modern terminology, anacidis implicitly a Brønsted acid and not a Lewis acid, since chemists almost always refer to a Lewis acid explicitly as such.

Definitions and concepts

Main article:Acid–base reaction

Modern definitions are concerned with the fundamental chemical reactions common to all acids.

Most acids encountered in everyday life areaqueous solutions,or can be dissolved in water, so the Arrhenius and Brønsted–Lowry definitions are the most relevant.

The Brønsted–Lowry definition is the most widely used definition; unless otherwise specified, acid–base reactions are assumed to involve the transfer of a proton (H⁺) from an acid to a base.

Hydronium ions are acids according to all three definitions. Although alcohols and amines can be Brønsted–Lowry acids, they can also function asLewis basesdue to the lone pairs of electrons on their oxygen and nitrogen atoms.

Arrhenius acids

Svante Arrhenius

In 1884,Svante Arrheniusattributed the properties of acidity tohydrogen ions(H⁺), later described asprotonsorhydrons.AnArrhenius acidis a substance that, when added to water, increases the concentration of H⁺ions in the water.Chemists often write H⁺(aq) and refer to thehydrogen ionwhen describing acid–base reactions but the free hydrogen nucleus, aproton,does not exist alone in water, it exists as thehydronium ion(H₃O⁺) or other forms (H₅O₂⁺,H₉O₄⁺). Thus, an Arrhenius acid can also be described as a substance that increases the concentration of hydronium ions when added to water. Examples include molecular substances such as hydrogen chloride and acetic acid.

An Arrheniusbase,on the other hand, is a substance that increases the concentration ofhydroxide(OH⁻) ions when dissolved in water. This decreases the concentration of hydronium because the ions react to form H₂O molecules:

H₃O⁺
_(aq)+ OH⁻
_(aq)⇌ H₂O_(liq)+ H₂O_(liq)

Due to this equilibrium, any increase in the concentration of hydronium is accompanied by a decrease in the concentration of hydroxide. Thus, an Arrhenius acid could also be said to be one that decreases hydroxide concentration, while an Arrhenius base increases it.

In an acidic solution, the concentration of hydronium ions is greater than 10⁻⁷molesper liter. Since pH is defined as the negative logarithm of the concentration of hydronium ions, acidic solutions thus have a pH of less than 7.

Brønsted–Lowry acids

Main article:Brønsted–Lowry acid–base theory

Acetic acid,aweak acid,donates a proton (hydrogen ion, highlighted in green) to water in an equilibrium reaction to give theacetateion and thehydroniumion. Red: oxygen, black: carbon, white: hydrogen.

While the Arrhenius concept is useful for describing many reactions, it is also quite limited in its scope. In 1923, chemistsJohannes Nicolaus BrønstedandThomas Martin Lowryindependently recognized that acid–base reactions involve the transfer of a proton. ABrønsted–Lowry acid(or simply Brønsted acid) is a species that donates a proton to a Brønsted–Lowry base.Brønsted–Lowry acid–base theory has several advantages over Arrhenius theory. Consider the following reactions ofacetic acid(CH₃COOH), theorganic acidthat gives vinegar its characteristic taste:

CH₃COOH + H₂O ⇌ CH₃COO⁻+ H₃O⁺

CH₃COOH + NH₃⇌ CH₃COO⁻+ NH+4

Both theories easily describe the first reaction: CH₃COOH acts as an Arrhenius acid because it acts as a source of H₃O⁺when dissolved in water, and it acts as a Brønsted acid by donating a proton to water. In the second example CH₃COOH undergoes the same transformation, in this case donating a proton to ammonia (NH₃), but does not relate to the Arrhenius definition of an acid because the reaction does not produce hydronium. Nevertheless, CH₃COOH is both an Arrhenius and a Brønsted–Lowry acid.

Brønsted–Lowry theory can be used to describe reactions ofmolecular compoundsin nonaqueous solution or the gas phase.Hydrogen chloride(HCl) and ammonia combine under several different conditions to formammonium chloride,NH₄Cl. In aqueous solution HCl behaves ashydrochloric acidand exists as hydronium and chloride ions. The following reactions illustrate the limitations of Arrhenius's definition:

1. H₃O⁺
   _(aq)+ Cl⁻
   _(aq)+ NH₃→ Cl⁻
   _(aq)+ NH⁺
   _4(aq)+ H₂O
2. HCl_(benzene)+ NH_3(benzene)→ NH₄Cl_(s)
3. HCl_(g)+ NH_3(g)→ NH₄Cl_(s)

As with the acetic acid reactions, both definitions work for the first example, where water is the solvent and hydronium ion is formed by the HCl solute. The next two reactions do not involve the formation of ions but are still proton-transfer reactions. In the second reaction hydrogen chloride and ammonia (dissolved inbenzene) react to form solid ammonium chloride in a benzene solvent and in the third gaseous HCl and NH₃combine to form the solid.

Lewis acids

Main article:Lewis acids and bases

A third, only marginally related concept was proposed in 1923 byGilbert N. Lewis,which includes reactions with acid–base characteristics that do not involve a proton transfer. ALewis acidis a species that accepts a pair of electrons from another species; in other words, it is an electron pair acceptor.Brønsted acid–base reactions are proton transfer reactions while Lewis acid–base reactions are electron pair transfers. Many Lewis acids are not Brønsted–Lowry acids. Contrast how the following reactions are described in terms of acid–base chemistry:

In the first reaction afluoride ion,F⁻,gives up anelectron pairtoboron trifluorideto form the producttetrafluoroborate.Fluoride "loses" a pair ofvalence electronsbecause the electrons shared in the B—F bond are located in the region of space between the two atomicnucleiand are therefore more distant from the fluoride nucleus than they are in the lone fluoride ion. BF₃is a Lewis acid because it accepts the electron pair from fluoride. This reaction cannot be described in terms of Brønsted theory because there is no proton transfer.

The second reaction can be described using either theory. A proton is transferred from an unspecified Brønsted acid to ammonia, a Brønsted base; alternatively, ammonia acts as a Lewis base and transfers a lone pair of electrons to form a bond with a hydrogen ion. The species that gains the electron pair is the Lewis acid; for example, the oxygen atom in H₃O⁺gains a pair of electrons when one of the H—O bonds is broken and the electrons shared in the bond become localized on oxygen.

Depending on the context, a Lewis acid may also be described as anoxidizeror anelectrophile.Organic Brønsted acids, such as acetic, citric, or oxalic acid, are not Lewis acids.They dissociate in water to produce a Lewis acid, H⁺,but at the same time, they also yield an equal amount of a Lewis base (acetate, citrate, or oxalate, respectively, for the acids mentioned). This article deals mostly with Brønsted acids rather than Lewis acids.

Dissociation and equilibrium

Reactions of acids are often generalized in the formHA ⇌ H⁺+ A⁻,where HA represents the acid and A⁻is theconjugate base.This reaction is referred to asprotolysis.The protonated form (HA) of an acid is also sometimes referred to as thefree acid.

Acid–base conjugate pairs differ by one proton, and can be interconverted by the addition or removal of a proton (protonationanddeprotonation,respectively). The acid can be the charged species and the conjugate base can be neutral in which case the generalized reaction scheme could be written asHA⁺⇌ H⁺+ A.In solution there exists anequilibriumbetween the acid and its conjugate base. Theequilibrium constantKis an expression of the equilibrium concentrations of the molecules or the ions in solution. Brackets indicate concentration, such that [H₂O] meansthe concentration of H₂O.Theacid dissociation constantK_ais generally used in the context of acid–base reactions. The numerical value ofK_ais equal to theproduct(multiplication) of the concentrations of the products divided by the concentration of the reactants, where the reactant is the acid (HA) and the products are the conjugate base and H⁺.

$K_{a}={\frac {{\ce {[H+] [A^{-}]}}}{{\ce {[HA]}}}}$

The stronger of two acids will have a higherK_athan the weaker acid; the ratio of hydrogen ions to acid will be higher for the stronger acid as the stronger acid has a greater tendency to lose its proton. Because the range of possible values forK_aspans many orders of magnitude, a more manageable constant, pK_ais more frequently used, where pK_a= −log₁₀K_a.Stronger acids have a smaller pK_athan weaker acids. Experimentally determined pK_aat 25 °C in aqueous solution are often quoted in textbooks and reference material.

Nomenclature

Arrhenius acids are named according to theiranions.In the classical naming system, the ionic suffix is dropped and replaced with a new suffix, according to the table following. The prefix "hydro-" is used when the acid is made up of just hydrogen and one other element. For example, HCl haschlorideas its anion, so the hydro- prefix is used, and the -ide suffix makes the name take the formhydrochloric acid.

Classical naming system:

Anion prefix	Anion suffix	Acid prefix	Acid suffix	Example
per	ate	per	ic acid	perchloric acid(HClO4)
				chloric acid(HClO3)
	ite		ous acid	chlorous acid(HClO2)
hypo	ite	hypo	ous acid	hypochlorous acid(HClO)
	ide	hydro	ic acid	hydrochloric acid(HCl)

In theIUPACnaming system, "aqueous" is simply added to the name of the ionic compound. Thus, for hydrogen chloride, as an acid solution, the IUPAC name is aqueous hydrogen chloride.

Acid strength

Main article:Acid strength

The strength of an acid refers to its ability or tendency to lose a proton. A strong acid is one that completely dissociates in water; in other words, onemoleof a strong acid HA dissolves in water yielding one mole of H⁺and one mole of the conjugate base, A⁻,and none of the protonated acid HA. In contrast, a weak acid only partially dissociates and at equilibrium both the acid and the conjugate base are in solution. Examples ofstrong acidsarehydrochloric acid(HCl),hydroiodic acid(HI),hydrobromic acid(HBr),perchloric acid(HClO₄),nitric acid(HNO₃) andsulfuric acid(H₂SO₄). In water each of these essentially ionizes 100%. The stronger an acid is, the more easily it loses a proton, H⁺.Two key factors that contribute to the ease of deprotonation are thepolarityof the H—A bond and the size of atom A, which determines the strength of the H—A bond. Acid strengths are also often discussed in terms of the stability of the conjugate base.

Stronger acids have a largeracid dissociation constant,K_aand a lower pK_athan weaker acids.

Sulfonic acids,which are organic oxyacids, are a class of strong acids. A common example istoluenesulfonic acid(tosylic acid). Unlike sulfuric acid itself, sulfonic acids can be solids. In fact,polystyrenefunctionalized into polystyrene sulfonate is a solid strongly acidic plastic that is filterable.

Superacidsare acids stronger than 100% sulfuric acid. Examples of superacids arefluoroantimonic acid,magic acidandperchloric acid.The strongest known acid ishelium hydride ion,with aproton affinityof 177.8kJ/mol.Superacids can permanently protonate water to give ionic, crystallinehydronium"salts". They can also quantitatively stabilizecarbocations.

WhileK_ameasures the strength of an acid compound, the strength of an aqueous acid solution is measured by pH, which is an indication of the concentration of hydronium in the solution. The pH of a simple solution of an acid compound in water is determined by the dilution of the compound and the compound'sK_a.

Lewis acid strength in non-aqueous solutions

Lewis acidshave been classified in theECW modeland it has been shown that there is no one order of acid strengths.The relative acceptor strength of Lewis acids toward a series of bases, versus other Lewis acids, can be illustrated byC-B plots.It has been shown that to define the order of Lewis acid strength at least two properties must be considered. For Pearson's qualitativeHSAB theorythe two properties arehardnessand strength while for Drago's quantitativeECW modelthe two properties are electrostatic and covalent.

Chemical characteristics

Monoprotic acids

See also:Acid dissociation constant § Monoprotic acids

Monoprotic acids, also known as monobasic acids, are those acids that are able to donate oneprotonper molecule during the process ofdissociation(sometimes called ionization) as shown below (symbolized by HA):

HA (aq) + H₂O (l) ⇌ H₃O⁺(aq) + A⁻(aq)K_a

Common examples of monoprotic acids inmineral acidsincludehydrochloric acid(HCl) andnitric acid(HNO₃). On the other hand, fororganic acidsthe term mainly indicates the presence of onecarboxylic acidgroup and sometimes these acids are known as monocarboxylic acid. Examples inorganic acidsincludeformic acid(HCOOH),acetic acid(CH₃COOH) andbenzoic acid(C₆H₅COOH).

Polyprotic acids

See also:Acid dissociation constant § Polyprotic acids

Polyprotic acids, also known as polybasic 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 (or dibasic) acid (two potential protons to donate), and triprotic (or tribasic) acid (three potential protons to donate). Some macromolecules such as proteins and nucleic acids can have a very large number of acidic protons.

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_a1and K_a2.

H₂A (aq) + H₂O (l) ⇌ H₃O⁺(aq) + HA⁻(aq)K_a1

HA⁻(aq) + H₂O (l) ⇌ H₃O⁺(aq) + A²⁻(aq)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 thebisulfateanion (HSO⁻
₄), for whichK_a1is very large; then it can donate a second proton to form thesulfateanion (SO²⁻
₄), wherein theK_a2is intermediate strength. The largeK_a1for the first dissociation makes sulfuric a strong acid. In a similar manner, the weak unstablecarbonic acid(H₂CO₃)can lose one proton to formbicarbonateanion(HCO⁻
₃)and lose a second to formcarbonateanion (CO²⁻
₃). BothK_avalues are small, butK_a1>K_a2.

A triprotic acid (H₃A) can undergo one, two, or three dissociations and has three dissociation constants, whereK_a1>K_a2>K_a3.

H₃A (aq) + H₂O (l) ⇌ H₃O⁺(aq) + H₂A⁻(aq)K_a1

H₂A⁻(aq) + H₂O (l) ⇌ H₃O⁺(aq) + HA²⁻(aq)K_a2

HA²⁻(aq) + H₂O (l) ⇌ H₃O⁺(aq) + A³⁻(aq)K_a3

Aninorganicexample of a triprotic acid is orthophosphoric acid (H₃PO₄), usually just calledphosphoric acid.All three protons can be successively lost to yield H₂PO⁻
₄,then HPO²⁻
₄,and finally PO³⁻
₄,the orthophosphate ion, usually just calledphosphate.Even though the positions of the three protons on the original phosphoric acid molecule are equivalent, the successiveK_avalues differ since it is energetically less favorable to lose a proton if the conjugate base is more negatively charged. Anorganicexample of a triprotic acid iscitric acid,which can successively lose three protons to finally form thecitrateion.

Although the subsequent loss of each hydrogen ion is less favorable, all of the conjugate bases are present in solution. The fractional concentration,α(alpha), for each species can be calculated. For example, a generic diprotic acid will generate 3 species in solution: H₂A, HA⁻,and A²⁻.The fractional concentrations can be calculated as below when given either the pH (which can be converted to the [H⁺]) or the concentrations of the acid with all its conjugate bases:

${\begin{aligned}\alpha _{{\ce {H2A}}}&={\frac {{\ce {[H+]^2}}}{{\ce {[H+]^2}}+[{\ce {H+}}]K_{1}+K_{1}K_{2}}}={\frac {{\ce {[H2A]}}}{{\ce {{[H2A]}}}+[HA^{-}]+[A^{2-}]}}\\\alpha _{{\ce {HA^-}}}&={\frac {[{\ce {H+}}]K_{1}}{{\ce {[H+]^2}}+[{\ce {H+}}]K_{1}+K_{1}K_{2}}}={\frac {{\ce {[HA^-]}}}{{\ce {[H2A]}}+{[HA^{-}]}+{[A^{2-}]}}}\\\alpha _{{\ce {A^{2-}}}}&={\frac {K_{1}K_{2}}{{\ce {[H+]^2}}+[{\ce {H+}}]K_{1}+K_{1}K_{2}}}={\frac {{\ce {[A^{2-}]}}}{{\ce {{[H2A]}}}+{[HA^{-}]}+{[A^{2-}]}}}\end{aligned}}$

A plot of these fractional concentrations against pH, for givenK₁andK₂,is known as aBjerrum plot.A pattern is observed in the above equations and can be expanded to the generaln-protic acid that has been deprotonatedi-times:

$\alpha _{{\ce {H}}_{n-i}A^{i-}}={{[{\ce {H+}}]^{n-i}\displaystyle \prod _{j=0}^{i}K_{j}} \over {\displaystyle \sum _{i=0}^{n}{\Big [}[{\ce {H+}}]^{n-i}\displaystyle \prod _{j=0}^{i}K_{j}}{\Big ]}}$

whereK₀= 1 and the other K-terms are the dissociation constants for the acid.

Neutralization

Hydrochloric acid(inbeaker) reacting withammoniafumes to produceammonium chloride(white smoke).

Neutralizationis the reaction between an acid and a base, producing asaltand neutralized base; for example,hydrochloric acidandsodium hydroxideformsodium chlorideand water:

HCl_(aq)+ NaOH_(aq)→ H₂O_(l)+ NaCl_(aq)

Neutralization is the basis oftitration,where apH indicatorshows 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 a base weaker than the acid results in a weakly acidic salt. An example is the weakly acidicammonium chloride,which is produced from the strong acidhydrogen chlorideand the weak baseammonia.Conversely, neutralizing a weak acid with a strong base gives a weakly basic salt (e.g.,sodium fluoridefromhydrogen fluorideandsodium hydroxide).

Weak acid–weak base equilibrium

Main article:Henderson–Hasselbalch equation

In order for a protonated acid to lose a proton, the pH of the system must rise above the pK_aof the 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.

Solutions of weak acids and salts of their conjugate bases formbuffer solutions.

Titration

Main article:Acid–base titration

To determine the concentration of an acid in an aqueous solution, an acid–base titration is commonly performed. A strong base solution with a known concentration, usually NaOH or KOH, is added to neutralize the acid solution according to the color change of the indicator with the amount of base added.The titration curve of an acid titrated by a base has two axes, with the base volume on the x-axis and the solution's pH value on the y-axis. The pH of the solution always goes up as the base is added to the solution.

Example: Diprotic acid

This is an ideal titration curve foralanine,a diprotic amino acid.Point 2 is the first equivalent point where the amount of NaOH added equals the amount of alanine in the original solution.

For each diprotic acid titration curve, from left to right, there are two midpoints, two equivalence points, and two buffer regions.

Equivalence points

Due to the successive dissociation processes, there are two equivalence points in the titration curve of a diprotic acid.The first equivalence point occurs when all first hydrogen ions from the first ionization are titrated.In other words, the amount of OH⁻added equals the original amount of H₂A at the first equivalence point. The second equivalence point occurs when all hydrogen ions are titrated. Therefore, the amount of OH⁻added equals twice the amount of H₂A at this time. For a weak diprotic acid titrated by a strong base, the second equivalence point must occur at pH above 7 due to the hydrolysis of the resulted salts in the solution.At either equivalence point, adding a drop of base will cause the steepest rise of the pH value in the system.

Buffer regions and midpoints

A titration curve for a diprotic acid contains two midpoints where pH=pK_a.Since there are two different K_avalues, the first midpoint occurs at pH=pK_a1and the second one occurs at pH=pK_a2.Each segment of the curve that contains a midpoint at its center is called the buffer region. Because the buffer regions consist of the acid and its conjugate base, it can resist pH changes when base is added until the next equivalent points.

Applications of acids

In industry

Acids are fundamental reagents in treating almost all processes in modern industry. Sulfuric acid, a diprotic acid, is the most widely used acid in industry, and is also the most-produced industrial chemical in the world. It is mainly used in producing fertilizer, detergent, batteries and dyes, as well as used in processing many products such like removing impurities.According to the statistics data in 2011, the annual production of sulfuric acid was around 200 million tonnes in the world.For example, phosphate minerals react with sulfuric acid to producephosphoric acidfor the production of phosphate fertilizers, andzincis produced by dissolving zinc oxide into sulfuric acid, purifying the solution and electrowinning.

In the chemical industry, acids react in neutralization reactions to produce salts. For example,nitric acidreacts withammoniato produceammonium nitrate,a fertilizer. Additionally,carboxylic acidscan beesterifiedwith alcohols, to produceesters.

Acids are often used to remove rust and other corrosion from metals in a process known aspickling.They may be used as an electrolyte in awet cell battery,such assulfuric acidin acar battery.

In food

Carbonated water (H₂CO₃aqueous solution) is commonly added to soft drinks to make them effervesce.

Tartaric acidis an important component of some commonly used foods like unripened mangoes and tamarind. Natural fruits and vegetables also contain acids.Citric acidis present in oranges, lemon and other citrus fruits.Oxalic acidis present in tomatoes, spinach, and especially incarambolaandrhubarb;rhubarb leaves and unripe carambolas are toxic because of high concentrations of oxalic acid.Ascorbic acid(Vitamin C) is an essential vitamin for the human body and is present in such foods as amla (Indian gooseberry), lemon, citrus fruits, and guava.

Many acids can be found in various kinds of food as additives, as they alter their taste and serve as preservatives.Phosphoric acid,for example, is a component ofcoladrinks.Acetic acidis used in day-to-day life as vinegar. Citric acid is used as a preservative in sauces and pickles.

Carbonic acidis one of the most common acid additives that are widely added insoft drinks.During the manufacturing process, CO₂is usually pressurized to dissolve in these drinks to generate carbonic acid. Carbonic acid is very unstable and tends to decompose into water and CO₂at room temperature and pressure. Therefore, when bottles or cans of these kinds of soft drinks are opened, the soft drinks fizz and effervesce as CO₂bubbles come out.

Certain acids are used as drugs.Acetylsalicylic acid(Aspirin) is used as a pain killer and for bringing down fevers.

In human bodies

Acids play important roles in the human body. The hydrochloric acid present in the stomach aids digestion by breaking down large and complex food molecules. Amino acids are required for synthesis of proteins required for growth and repair of body tissues. Fatty acids are also required for growth and repair of body tissues. Nucleic acids are important for the manufacturing of DNA and RNA and transmitting of traits to offspring through genes. Carbonic acid is important for maintenance of pH equilibrium in the body.

Human bodies contain a variety of organic and inorganic compounds, among thosedicarboxylic acidsplay an essential role in many biological behaviors. Many of those acids areamino acids,which mainly serve as materials for the synthesis of proteins.Other weak acids serve as buffers with their conjugate bases to keep the body's pH from undergoing large scale changes that would be harmful to cells.The rest of the dicarboxylic acids also participate in the synthesis of various biologically important compounds in human bodies.

Acid catalysis

Main article:Acid catalysis

Acids are used ascatalystsin industrial and organic chemistry; for example,sulfuric acidis used in very large quantities in thealkylationprocess to produce gasoline. Some acids, such as sulfuric, phosphoric, and hydrochloric acids, also effectdehydrationandcondensation reactions.In biochemistry, manyenzymesemploy acid catalysis.

Biological occurrence

Basic structure of anamino acid.

Many biologically important molecules are acids.Nucleic acids,which contain acidicphosphate groups,includeDNAandRNA.Nucleic acids contain the genetic code that determines many of an organism's characteristics, and is passed from parents to offspring. DNA contains the chemical blueprint for the synthesis ofproteins,which are made up ofamino acidsubunits.Cell membranescontainfatty acidesterssuch asphospholipids.

An α-amino acid has a central carbon (the α oralphacarbon) that is covalently bonded to acarboxylgroup (thus they arecarboxylic acids), anaminogroup, a hydrogen atom and a variable group. The variable group, also called the R group or side chain, determines the identity and many of the properties of a specific amino acid. Inglycine,the simplest amino acid, the R group is a hydrogen atom, but in all other amino acids it is contains one or more carbon atoms bonded to hydrogens, and may contain other elements such as sulfur, oxygen or nitrogen. With the exception of glycine, naturally occurring amino acids arechiraland almost invariably occur in theL-configuration.Peptidoglycan,found in some bacterialcell wallscontains someD-amino acids. At physiological pH, typically around 7, free amino acids exist in a charged form, where the acidic carboxyl group (-COOH) loses a proton (-COO⁻) and the basic amine group (-NH₂) gains a proton (-NH⁺
₃). The entire molecule has a net neutral charge and is azwitterion,with the exception of amino acids with basic or acidic side chains.Aspartic acid,for example, possesses one protonated amine and two deprotonated carboxyl groups, for a net charge of −1 at physiological pH.

Fatty acids and fatty acid derivatives are another group of carboxylic acids that play a significant role in biology. These contain long hydrocarbon chains and a carboxylic acid group on one end. The cell membrane of nearly all organisms is primarily made up of aphospholipid bilayer,amicelleof hydrophobic fatty acid esters with polar, hydrophilicphosphate"head" groups. Membranes contain additional components, some of which can participate in acid–base reactions.

In humans and many other animals,hydrochloric acidis a part of thegastric acidsecreted within thestomachto help hydrolyzeproteinsandpolysaccharides,as well as converting the inactive pro-enzyme,pepsinogeninto theenzyme,pepsin.Some organisms produce acids for defense; for example, ants produceformic acid.

Acid–base equilibrium plays a critical role in regulatingmammalianbreathing.Oxygengas (O₂) drivescellular respiration,the process by which animals release the chemicalpotential energystored in food, producingcarbon dioxide(CO₂) as a byproduct. Oxygen and carbon dioxide are exchanged in thelungs,and the body responds to changing energy demands by adjusting the rate ofventilation.For example, during periods of exertion the body rapidly breaks down storedcarbohydratesand fat, releasing CO₂into the blood stream. In aqueous solutions such as blood CO₂exists in equilibrium withcarbonic acidandbicarbonateion.

CO₂+ H₂O ⇌ H₂CO₃⇌ H⁺+ HCO−3

It is the decrease in pH that signals the brain to breathe faster and deeper, expelling the excess CO₂and resupplying the cells with O₂.

Aspirin(acetylsalicylic acid) is acarboxylic acid

Cell membranesare generally impermeable to charged or large, polar molecules because of thelipophilicfatty acyl chains comprising their interior. Many biologically important molecules, including a number of pharmaceutical agents, are organic weak acids that can cross the membrane in their protonated, uncharged form but not in their charged form (i.e., as the conjugate base). For this reason the activity of many drugs can be enhanced or inhibited by the use of antacids or acidic foods. The charged form, however, is often more soluble in blood andcytosol,both aqueous environments. When the extracellular environment is more acidic than the neutral pH within the cell, certain acids will exist in their neutral form and will be membrane soluble, allowing them to cross the phospholipid bilayer. Acids that lose a proton at theintracellular pHwill exist in their soluble, charged form and are thus able to diffuse through the cytosol to their target.Ibuprofen,aspirinandpenicillinare examples of drugs that are weak acids.

Common acids

Mineral acids (inorganic acids)

• Hydrogen halidesand their solutions:hydrofluoric acid(HF),hydrochloric acid(HCl),hydrobromic acid(HBr),hydroiodic acid(HI)
• Halogen oxoacids:hypochlorous acid(HClO),chlorous acid(HClO₂),chloric acid(HClO₃),perchloric acid(HClO₄), and corresponding analogs for bromine and iodine
  • Hypofluorous acid(HFO), the only known oxoacid for fluorine.
• Sulfuric acid(H₂SO₄)
• Fluorosulfuric acid(HSO₃F)
• Nitric acid(HNO₃)
• Phosphoric acid(H₃PO₄)
• Fluoroantimonic acid(HSbF₆)
• Fluoroboric acid(HBF₄)
• Hexafluorophosphoric acid(HPF₆)
• Chromic acid(H₂CrO₄)
• Boric acid(H₃BO₃)

Sulfonic acids

Asulfonic acidhas the general formula RS(=O)₂–OH, where R is an organic radical.

• Methanesulfonic acid(or mesylic acid, CH₃SO₃H)
• Ethanesulfonic acid(or esylic acid, CH₃CH₂SO₃H)
• Benzenesulfonic acid(or besylic acid, C₆H₅SO₃H)
• p-Toluenesulfonic acid(or tosylic acid, CH₃C₆H₄SO₃H)
• Trifluoromethanesulfonic acid(or triflic acid, CF₃SO₃H)
• Polystyrene sulfonic acid(sulfonatedpolystyrene,[CH₂CH(C₆H₄)SO₃H]_n)

Carboxylic acids

Acarboxylic acidhas the general formula R-C(O)OH, where R is an organic radical. The carboxyl group -C(O)OH contains acarbonylgroup, C=O, and ahydroxylgroup, O-H.

• Acetic acid(CH₃COOH)
• Citric acid(C₆H₈O₇)
• Formic acid(HCOOH)
• Gluconic acidHOCH₂-(CHOH)₄-COOH
• Lactic acid(CH₃-CHOH-COOH)
• Oxalic acid(HOOC-COOH)
• Tartaric acid(HOOC-CHOH-CHOH-COOH)

Halogenated carboxylic acids

Halogenation atalpha positionincreases acid strength, so that the following acids are all stronger than acetic acid.

• Fluoroacetic acid
• Trifluoroacetic acid
• Chloroacetic acid
• Dichloroacetic acid
• Trichloroacetic acid

Vinylogous carboxylic acids

Normal carboxylic acids are the direct union of a carbonyl group and a hydroxyl group. Invinylogouscarboxylic acids, a carbon-carbon double bond separates the carbonyl and hydroxyl groups.

• Ascorbic acid

Nucleic acids

• Deoxyribonucleic acid(DNA)
• Ribonucleic acid(RNA)

References

• Listing of strengths of common acids and bases
• Zumdahl, Steven S. (1997).Chemistry(4th ed.). Boston: Houghton Mifflin.ISBN9780669417944.
• Pavia, D. L.; Lampman, G. M.; Kriz, G. S. (2004).Organic Chemistry Volume I.Mason, OH: Cengage Learning.ISBN0759347271.

External links

• Curtipot:Acid–Base equilibria diagrams,pHcalculation andtitrationcurves simulation and analysis –freeware