Acid

For other senses of this word, see acid (disambiguation).

Acids and bases:

• Acid-base extraction
• Acid-base reaction
• Acid dissociation constant
• Acidity function
• Buffer solutions
• pH
• Proton affinity
• Self-ionization of water
• Acids:
  • Lewis acids
  • Mineral acids
  • Organic acids
  • Strong acids
  • Superacids
  • Weak acids
• Bases:
  • Lewis bases
  • Organic bases
  • Strong bases
  • Superbases
  • Non-nucleophilic bases
  • Weak bases

An acid (often represented by the generic formula HA) is traditionally considered any chemical compound that, when dissolved in water, gives a solution with a pH of less than 7.0. That approximates the modern definition of Brønsted and Lowry, who defined an acid as a compound which donates a hydrogen ion (H⁺) to another compound (called a base). Common examples include acetic acid (in vinegar) and sulfuric acid (used in car batteries).

Definitions of acids and bases

The word "acid" comes from the Latin acidus meaning "sour," but in chemistry the term acid has a more specific meaning. There are three common ways to define an acid, namely, the Arrhenius, the Brønsted-Lowry and the Lewis definitions, in order of increasing generality.

• Arrhenius: According to this definition, an acid is a substance that increases the concentration of hydronium ion (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 French chemists, including Antoine Lavoisier, incorrectly believed that all acids contained oxygen. Indeed the modern German word for Oxygen is Sauerstoff (lit. sour substance). English chemists, including Sir Humphry Davy at the same time believed all acids contained hydrogen. The Swedish chemist Svante 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 (hydrogen nucleus) 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 formulated this definition, which includes water-insoluble substances not in the Arrhenius definition.
• Lewis: According to this definition, an acid is an electron-pair acceptor and a base is an electron-pair donor. (These are frequently referred to as "Lewis acids" and "Lewis bases," and are electrophiles and nucleophiles, respectively, in organic chemistry; Lewis bases are also ligands in 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. 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. This definition was developed by Gilbert N. Lewis.

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 stability of the conjugate base will increase the acidity of a compound. This concept of acidity is used frequently for organic acids such as carboxylic acid. The molecular orbital description, where the unfilled proton orbital overlaps with a lone pair, is connected to the Lewis definition.

Solutions of weak acids and salts of their conjugate bases form buffer solutions.

Acid/base systems are different from redox reactions in that there is no change in oxidation state.

Properties

Generally, acids have the following properties:

• Taste: Acids generally are sour when dissolved in water.
• Touch: Acids produce a stinging feeling, particularly strong acids.
• Reactivity: Acids react aggressively with or corrode most metals.
• Electrical conductivity: Acids, while not normally ionic, are electrolytes.

Strong acids and most concentrated acids are dangerous, causing severe burns for even minor contact. 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 reaction of the acid dissolving in the water could cause severe thermal burns. In addition to dangers from the acidity, even dilute solutions of weak acids may also be dangerous, due to toxic or other effects of the ions involved. See an appropriate MSDS for more specific information.

Nomenclature

Acids are named according to their anions. 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.

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 equilibrium occurs between an acid (HA) and water, which acts as a base:

HA(aq) ⇌ H₃O⁺(aq) + A^-(aq)

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

$K_{a}={[{\mbox{H}}_{3}{\mbox{O}}^{+}]\cdot [A^{-}] \over [HA]}$

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 acids: 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 acids 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. Oxoacids, 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 pK_a (pK_a= - log(K_a)).
• The pH of a solution is a measurement of the concentration of hydronium. This will depend of 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 (sometimes called ionization) as follows and simply has one acid dissociation constant as shown above:

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

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 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 and 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.

Weak acid/weak base equilibria

In order to lose a proton, it is necessary that the pH of the system rise above the pK_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).

Acidification of the environment

Acidification is the process whereby a compound is added to a solution, leading to a drop in the pH of the solution. One example is when the pollution of air – mainly sulfur dioxide and nitrogen oxides – is converted into acidic substances.

This ‘acid rain’ is best known for the damage it causes to forests and lakes. It also damages freshwater and coastal ecosystems, soils and even ancient historical monuments, or the heavy metals these acids help release into groundwater.

Sulfur dioxide and the nitrogen oxides are mainly emitted by burning fossil fuels. The 1990s saw these emissions drop substantially, thanks to a combination of European Directives forcing the installation of desulfurisation systems, the move away from coal as a fossil fuel, and major economic restructuring in the new German Lander.

Acidification is nevertheless still a major environmental problem in Europe. It is a cross-border issue, requiring coordinated initiatives across countries and sectors. This section brings together the EEA’s reports on the scale of the problem and the effectiveness of the solutions tried to date.

See also

Chemistry

• acid number
• acid salt
• Base (chemistry)
• basic salt
• vitriol

Environment

• acid rain
• ocean acidification

References

ISBN links support NWE through referral fees

• Listing of strengths of common acids and bases
• Zumdahl, Chemistry 4th Edition.

External links

• Curtipot: Acid-Base equilibria diagrams, pH calculation and titration curves simulation and analysis - freeware

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