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Acid-base reaction theories
From Wikipedia, the free encyclopedia
- Acid-base reaction theories
- pH
- Self-ionization of water
- Buffer solutions
- Systematic naming
- Electrochemistry
- Acid-base extraction
- Acids:
- Bases:
An acid-base reaction is a chemical reaction that occurs between an acid and a base. Several concepts exist which provide alternative definitions for the reaction mechanisms involved and their application in solving related problems. Despite several similarities in definitions, their importance becomes apparent as different methods of analysis when applied to acid-base reactions for gaseous or liquid species, or when acid or base character may be somewhat less apparent.
Historically, the first of these scientific concepts of acids and bases was provided by the French chemist Antoine Lavoisier, circa 1776.[1]
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Since Lavoisier's knowledge of strong acids was mainly restricted to 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" (from the Greek οξυς (oxys) (acid, sharp) and γεινομαι (geinomai) (engender)). 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 of acid-base reactions is a more simplified acid-base concept devised by Svante Arrhenius, which was used to provide a modern definition of bases that followed from his work with Friedrich Wilhelm Ostwald in establishing the presence of ions in aqueous solution in 1884, and led to Arrhenius receiving the Nobel prize in chemistry in 1903.[2]
As defined at the time of discovery, acid-base reactions are characterized by Arrhenius acids, which dissociate in aqueous solution form hydrogen or oxonium (H3O) ions[2], and Arrhenius bases which form hydroxide (OH − ) ions. More succinctly;
| “ | Arrhenius acids form hydrogen ions in aqueous solution with Arrhenius bases forming hydroxide ions | ” |
The universal aqueous acid-base definition of the Arrhenius concept is described as the formation of water from hydrogen and hydroxide ions, or oxonium ions and hydroxide ions produced from the dissociation of an acid and base in aqueous solution (
)[3], which leads to the definition that in Arrhenius acid-base reactions, a salt and water is formed from the reaction between an acid and a base.[2]
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 sulphuric acid (H2SO4) to form two moles of water and one mole of sodium sulphate.
The Brønsted-Lowry definition, formulated independently by its two proponents Johannes Nicolaus Brønsted and Martin Lowry in 1923 is based upon the idea of protonation of bases through the de-protonation of acids -- more commonly referred to as the ability of acids to "donate" hydrogen ions (H + ) or protons to bases, which "accept" them.[4] In contrast to the Arrhenius definition, the Brønsted-Lowry definition refers to the products of an acid-base reaction as conjugate acids and bases to refer to the relation of one proton, and to indicate that there has been a reaction between the two quantities, rather than a "formation" of salt and water, as explained in the Arrhenius definition.[2][4]
It defines that in reactions, there is the donation and reception of a proton, which essentially refers to the removal of a hydrogen ion bonded within a compound and it's reaction with another compound[5], and not the removal of a proton from the nucleus of an atom, which would require inordinate amounts of energy not attainable through the simple dissociation of acids. In differentiation from the Arrhenius definition, the Brønsted-Lowry definition postulates that for each acid, there is a conjugate base that is formed through a complete reaction, which also includes water, which is amphoteric[5][2][4] :
-
- Hydrochloric acid completely reacts with water to form the hydronium and chlorine ions
-
- Acetic acid reacts incompletely with ammonia, no hydronium ions being produced
-
- General formula for representing Brønsted-Lowry reactions.
This definition is based on a generalization of the earlier Arrhenius definition to all autodissociating solvents. In all such solvents there is a certain concentration of a positive species, solvonium cations and negative species, solvate anions, in equilibrium with the neutral solvent molecules. For example:
or even some aprotic systems
- N2O4 ⇌ NO+ (nitrosonium) + NO3− (nitrate)
- 2SbCl3 ⇌ SbCl2+ (dichloroantimonium) + SbCl4- (tetrachloroantimonate)
A solute causing an increase in the concentration of the solvonium ions and a decrease in the solvate ions is an acid and one causing the reverse is a base. Thus, in liquid ammonia, KNH2 (supplying NH2-) is a strong base, and NH4NO3 (supplying NH4+) is a strong acid. In liquid sulfur dioxide (SO2), thionyl compounds (supplying SO2+) behave as acids, and sulfites (supplying SO32−) behave as bases.
Here are some nonaqueous acid-base reactions in liquid ammonia
- 2NaNH2 (base) + Zn(NH2)2 (amphiphilic amide) = Na2[Zn(NH2)4]
- 2NH4I (acid) + Zn(NH2)2 (amphiphilic amide) = [Zn(NH3)4)]I2
Nitric acid can be a base in liquid sulphuric acid:
- HNO3 (base) + 2H2SO4 = NO2+ + H3O+ + 2HSO4-
And things become even stranger in the aprotic world, for example in liquid N2O4:
- AgNO3 (base) + NOCl (acid) = N2O4 + AgCl
Since solvent-system definition depends on the solvent as well as on the compound itself, the same compound can change its role depending on the choice of the solvent. Thus, HClO4 is a strong acid in water, a weak acid in acetic acid, and a weak base in fluorosulfonic acid.
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. Such a covalent bond, in which both of the shared electrons originate from one of the reacting molecules, is known as a coordinate covalent bond. From the perspective of Molecular Orbital theory, 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 broadest acid-base definitions and is necessary for an understanding of many acid-base reactions. Even so, the Brønsted-Lowry definition is sufficient and more practical in most cases for everyday use.[citation needed]
See also Lewis acid and Lewis base.
The most general definition is that of the Russian chemist Mikhail 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.
- ^ Miessler, L. M., Tar, D. A., (1991) p166 - Table of discoveries attributes Antoine Lavoisier as the first to posit a scientific theory in relation to oxyacids.
- ^ a b c d e f Miessler, L. M., Tar, D. A., (1991) p165
- ^ The use of H + is as shorthand for H3O;
- ^ a b c Miessler, L. M., Tar, D. A., (1991), p167-169 -- According to this page, the original definition was that "acids have a tendency to lose a proton"
- ^ a b Clayden, J., Warren, S., et al. (2000), p182-184
- Miessler, L. M., Tar, D. A., (1991) "Inorganic Chemistry" 2nd ed. Pearson Prentice-Hall
- Clayden, J., Warren, S., et al. (2000) "Organic Chemistry" Oxford University Press