Acid-base reaction theories
|
|
| Acids and Bases: |
|---|
| Acid-base reaction theories |
| pH |
| Self-ionization of water |
| Buffer solutions |
| Systematic naming |
| Redox reactions |
| Electrochemistry |
| Strong acids |
| Weak acids |
| Strong bases |
| Weak bases |
An acid-base reaction is a chemical reaction between an acid and a base.
| Contents |
Common acid-base theories
Lavoisier's definition
The first scientific definition was proposed by the French chemist Antoine Lavoisier.
Since Lavoisier's knowledge of strong acids was mainly restricted to the oxyacids, which tend to contain central atoms in high oxidation states surrounded by oxygen, such as HNO3 and H2SO4, and since he was not aware of the true composition of the hydrohalic acids, HCl, HBr, and HI, he defined acids in terms of their containing oxygen, which in fact he named from Greek words meaning "acid-former". When the elements chlorine, bromine, and iodine were identified and the absence of oxygen in the hydrohalic acids was established by Sir Humphry Davy in 1810, this definition had to be rejected.
The Arrhenius definition
Svante Arrhenius provided the first modern definition of acids and bases in 1884. In water, a dissociation takes place:
- H2O → H+ + OH-
A compound causing an increase in H+ and a decrease in OH- is an acid and one causing the reverse is a base.
An Arrhenius acid, when dissociated in water, typically yields positively-charged hydrogen ion and a complementary negative ion.
An Arrhenius base, when dissociated in water, typically yields negatively-charged hydroxide ion and a complementary positive ion.
The positive ion from a base can form a salt with the negative ion from an acid. For example, two moles of the base sodium hydroxide (NaOH) can combine with one mole of sulfuric acid (H2SO4) to form two moles of water and one mole of sodium sulfate.
- 2NaOH + H2SO4 → 2H2O + Na2SO4
The protonic (Brønsted-Lowry) definition
The Brønsted-Lowry definition, formulated independently by its two proponents Johannes Nicolaus Brønsted and Martin Lowry in 1923, revolves around an acid's ability to donate protons (H+) to another compound, called a base, in a chemical reaction.
A base is a proton acceptor. In Brønsted-Lowry acid-base reactions, there is a "competition" between two bases for a proton. so that if X and Y are two species, the equilibrium
- HX + Y- ↔ HY + X-
occurs. Both HX and HY are Brønsted-Lowry acids; both X- and Y- are Brønsted-Lowry bases. If the reaction runs mostly to the left, then HY is the stronger acid and X- the stronger base; if the reaction runs mostly to the right, then HX is the stronger acid and Y- the stronger base.
It may be more intuitive to define the stronger of two acids as the one which reacts more completely with a common base. The following shows that this definition gives the same result. Compare the reactions of the two acids HX and HY with the same base Z- (in a mixture containing all these species):
- HX + Z- ↔ HZ + X-
- HY + Z- ↔ HZ + Y-
If these reactions have equilibrium constants KX and KY respectively, then:
- [X-][HZ] / [HX][Z-]=KX
- [Y-][HZ] / [HY][Z-]=KY
and hence (dividing):
- [X-][HY] / [HX][Y-] = KX / KY
Given that this last quantity is the equilibrium constant for the above reaction, the reaction will tend to the right if KX / KY > 1, in other words if HX is a stronger acid than HY under this definition, and vice versa.
Acids and bases in the Brønsted-Lowry system occur in conjugate pairs; in the reaction
- HX → H+ + X-
HX is called the conjugate acid of the base X-, and X- is called the conjugate base of the acid HX.
Some compounds, like water, can act either as an acid or a base, and are called amphoteric compounds.
Stronger acids typically oxidize metals, forming salts and releasing hydrogen.
See pH for a measure of proton concentration frequently used for measuring acidity and alkalinity using this definition. For a list of conjugate acid-base pairs, see conjugate acid.
The solvent-system definition
This definition is based on a generalization of the earlier Arrhenius definition. If we consider a solvent which can be dissociated into a positive species X and a negative species Y:
- XY ↔ X+ + Y-
or
- 2XY ↔ X2Y+ + Y-
or
- 2XY ↔ X+ + XY2-
a compound causing an increase in X+ (or X2Y+) and a decrease in Y- (or XY2-) is an acid and one causing the reverse is a base. For example in liquid sulfur dioxide (SO2), thionyl compounds (formally supplying SO2+) behave as acids, and sulfites (supplying SO32-) behave as bases.
In this more general sense, aprotic compounds (those which do not donate protons), can still react with bases, and the terms "acid" and "base" can still be used for reactions in aprotic or non-aqueous environments.
The electronic (Lewis) definition
The more general definition offered by Lewis in 1923 (the same year as the Brønsted-Lowry definition) describes the reactivity of an acid in terms of its ability to accept a pair of electrons from a base, defined as an electron-pair donor. In general, an acid reacts with a base by forming a new covalent bond utilizing an empty orbital of the acid to share the extra electron pair of the base. That is an acid-base reaction is the combination of HOMO from base and LUMO from acid to form a stable bonding molecular orbital.
The Lewis definition is one of the most broad definitions and is necessary for an understanding of acid-base reaction, although the Brønsted-Lowry definition is sufficient and more practical for most cases in everyday use.
Other acid-base theories
The Usanovich definition
The most general definition is that of the Russian chemist Usanovich, and can basically be summarized as defining an acid as anything that accepts negative species or donates positive ones, and a base as the reverse. This tends to overlap the concept of redox (oxidation-reduction), and so is not highly favored by chemists. This is because redox reactions focus more on physical electron transfer processes, rather than bond making/bond breaking processes, although the distinction between these two processes is somewhat ambiguous.