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 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). Common examples include acetic acid (in vinegar) and sulfuric acid (used in car
batteries). Acid/base systems are different from redox reactions in that there is no change
in oxidation state. Generally, acids have the following properties:
- Taste:Acids often taste sour
- Touch: Strong or concentrated acids often produce a stinging feeling on mucous
membranes
- Reactivity: Strong acids react aggressively with or corrode many metals
- Electrical conductivity: Acids, while not usually ionic compounds, are
electrolytes
- Acids turn moist blue litmus paper red
Definitions
-
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 Swedish chemist
Svante Arrhenius, 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), as is the Afrikaans word for oxygen suurstof, with the same
meaning. 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 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.
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 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.
Properties
Strong acids and many concentrated acids are dangerous, causing severe burns for even
minor contact. They are said to be 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 specific information.
Bronsted-Lowry Acids:
- Are generally sour in taste
- Turn blue litmus red
- Turn methyl orange red
- Do not change the colour of a solution of phenolphthalein, a common pH indicator (remains colourless)
- Will react with metals to produce a 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 only in aqueous solutions
- Will produce hydronium (H3O+) ions when dissolved in aqueous media
- Will denature proteins
Nomenclature
In the classical naming system, 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. 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:
Chemical characteristics
In water the following equilibrium occurs between a weak acid (HA) and water,
which acts as a base:
HA(aq) + H2O ⇌
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[\mbox{A}^-] \over [\mbox{HA}]}](http://content.answers.com/main/content/wp/en/math/d/2/a/d2a64114d433d8139a65d414e0b7d37e.png)
Strong acids have large Ka 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 Ka value for hydrochloric acid (HCl) is 107.
Weak acids have small Ka 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 Ka 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
(Ka) or equivalently its pKa (pKa= - log(Ka)).
- 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 (sometimes called ionization)
as follows and simply has one acid dissociation constant as shown above:
-
-
-
-
- HA(aq) + H2O(l) ⇌ H3O+(aq) + A−(aq)
Ka
A diprotic acid (here symbolized by H2A) can undergo one or two dissociations depending on the pH. Each
dissociation has its own dissociation constant, Ka1 and Ka2.
-
-
-
-
- H2A(aq) + H2O(l) ⇌ H3O+(aq) + HA−(aq)
Ka1
-
-
-
-
- HA−(aq) + H2O(l) ⇌ H3O+(aq) +
A2−(aq) Ka2
The first dissociation constant is typically greater than the second; i.e., Ka1 > Ka2 .
For example, sulfuric acid (H2SO4) can donate one proton to form the
bisulfate anion (HSO4−), for which Ka1 is very large;
then it can donate a second proton to form the sulfate anion (SO42−),
wherein the Ka2 is intermediate strength. The large Ka1 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 Ka values are small, but Ka1 > Ka2
.
A triprotic acid (H3A) can undergo one, two, or three dissociations and has three dissociation constants, where
Ka1 > Ka2 > Ka3 .
-
-
-
-
- H3A(aq) + H2O(l) ⇌ H3O+(aq) +
H2A−(aq) Ka1
-
-
-
-
- H2A−(aq) + H2O(l) ⇌ H3O+(aq) +
HA2−(aq) Ka2
-
-
-
-
- HA2−(aq) + H2O(l) ⇌ H3O+(aq) + A3−(aq)
Ka3
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 Ka 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. 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.
Weak acid/weak base equilibria
-
In order to lose a proton, it is necessary that the pH of the system rise above the pKa 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 solutions.
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. 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
proteins and polysaccharides, as well as converting the
inactive pro-enzyme, pepsinogen into the enzyme, pepsin. Acids
are used as catalysts; for example, sulfuric acid is
used in very large quantities in the alkylation process to produce gasoline.
Common Acids
Mineral Acids
(The following three are also known as the bench acids)
Other acids include:
(Misc)
(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))
References
See also
- Chemistry
- Environment
External links
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