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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 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) Ì! H3O+(aq) + A-(aq)
The acidity constant (or acid dissociation constant) is the equilibrium constant for the
reaction of HA with water:
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-). 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.
Strong acids include the hydrohalic acids - HCl, HBr, and HI. (However, hydrofluoric
acid, HF, is relatively weak.) Oxoacids, which tend to contain central atoms in high
oxidation states surrounded by oxygen, are also quite strong and include HNO3, H2SO4,
HClO4. Most organic acids are weak acids.
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 (Ka) or
equivalently its pKa (pKa= - log(Ka).
" The pH of a solution is a measurement of the concentration of hydronium. This
will depend of the concnetration 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 potenital protons to donate)
and triprotic acid (three potenital 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) Ì! 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) Ì! H3O+(aq) + HA-(aq) Ka1
HA-(aq) + H2O(l) Ì! 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) Ì! H3O+(aq) + H2A-(aq) Ka1
H2A-(aq) + H2O(l) Ì! H3O+(aq) + HA2-(aq) Ka2
HA2-(aq) + H2O(l) Ì! 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 [ Pobierz całość w formacie PDF ]