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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. 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). Acids generally taste sour; however, tasting acids, particularly concentrated acids, can be dangerous and is not recommended.

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. 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 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 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 in organic chemistry). Lewis acids include substances with no protons, such as iron(III) chloride. 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 defintion 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.

Generally, acids have the following chemical and physical 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 are electrolytes.

Strong acids are dangerous, causing severe burns for even minor contact. Generally, acid burns are treated by rinsing the affected area abundantly with water and followed up with immediate medical attention.

Nomenclature

Acids are named according to the ending of their anion. That ionic ending is dropped and replaced with a new suffix 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 Ending	Acid Prefix	Acid Suffix
per-anion-ate	per	ic acid
ate		ic acid
ite		ous acid
hypo-anion-ite	hypo	ous acid
ide	Hydro	ic acid

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 the following:

The terms "hydrogen ion" and "proton" are used interchangebly; 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.

Neutralisation

Neutralisation is the reaction between equal amounts of 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)

Neutralisation 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.

Common acids

Strong inorganic acids

Medium to weak inorganic acids

Weak organic acids

Acids in food

Acetic acid: (E260) found in vinegar
Adipic acid: (E355)
Alginic acid: (E400)
Ascorbic acid (vitamin C): (E300) found in fruits
Benzoic acid: (E210)
Boric acid: (E284)
Citric acid: (E330) found in citrus fruits
Carbonic acid: (E290) found in carbonated soft drinks
Carminic acid: (E120)
Cyclamic acid: (E952)
Erythorbic acid: (E315)
Erythorbin acid: (E317)
Formic acid: (E236)found in bee and ant stings
Fumaric acid: (E297)
Gluconic acid: (E574)
Glutamic acid: (E620)
Guanylic acid: (E626)
Hydrochloric acid: (E507)
Inosinic acid: (E630)
Lactic acid: (E270) found in dairy products such as yoghurt and sour milk, also is product of cellular fermentation, the reason muscles burn
Malic acid: (E296)
Metatartaric acid: (E353)
Methanethiol: found in cheese and some other fermented foods.
Niacin (nicotinic acid): (E375)
Oxalic acid: found in spinach and rhubarb
Pectic acid: found in fruits and some vegetables
Phosphoric acid: (E338)
Propionic acid: (E280)
Sorbic acid: (E200) found in foods and drinks
Stearic acid: (E570), a type of fatty acid.
Succinic acid: (E363)
Sulfuric acid: (E513)
Tannic acid: found in tea
Tartaric acid: (E334) found in grapes

Sources

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

@@ Line 97: / Line 97: @@
 An [[inorganic]] example of a triprotic acid is orthophosphoric acid (H<sub>3</sub>PO<sub>4</sub>), usually just called [[phosphoric acid]]. All three protons can be successively lost to yield H<sub>2</sub>PO<sub>4</sub><sup>&minus;</sup>, then HPO<sub>4</sub><sup>2&minus;</sup>, and finally PO<sub>4</sub><sup>3&minus;</sup> , 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''<sub>a</sub> values will differ since it is energetically less favorable to lose a proton if the conjugate base is more negatively charged.
-== Neutralization ==
+== Neutralisation ==
-[[Neutralization]] is the reaction between equal amounts of an acid and a base, producing a [[salt]] and [[water (molecule)|water]]; for example, hydrochloric acid and sodium hydroxide form sodium chloride and water:
+[[Neutralisation]] is the reaction between equal amounts of an acid and a base, producing a [[salt]] and [[water (molecule)|water]]; for example, hydrochloric acid and sodium hydroxide form sodium chloride and water:
 ::HCl(aq) + NaOH(aq) -> H<sub>2</sub>O(l) + NaCl(aq)
-Neutralization is the basis of [[titration]], where a [[PH indicator|pH indicator]] shows equivalence point when the equivalent number of moles of a base have been added to an acid.
+Neutralisation is the basis of [[titration]], where a [[PH indicator|pH indicator]] shows equivalence point when the equivalent number of moles of a base have been added to an acid.
 == Common acids ==