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An acid-base reaction is a chemical reaction that occurs between an acid and a base. Several concepts that provide alternative definitions for the reaction mechanisms involved and their application in solving related problems exist. Despite several differences 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. The first of these scientific concepts of acids and bases was provided by the French chemist Antoine Lavoisier, circa 1776.[1]
[edit] Common acid-base theories[edit] Lavoisier definitionSince Lavoisier's knowledge of strong acids was mainly restricted to oxoacids, 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 (HF, 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) meaning "acid" or "sharp" and γεινομαι (geinomai) meaning "engender"). The Lavoisier definition was held as absolute truth for over 30 years, until the 1810 article and subsequent lectures by Sir Humphry Davy in which he proved the lack of oxygen in H2S, H2Te, and the hydrohalic acids. [edit] Liebig definitionThis definition is proposed by Justus von Liebig circa 1838,[2] based on his extensive works on the chemical composition of organic acids. This finished the doctrinal shift from oxygen-based acids to hydrogen-based acids, started by Davy. According to Liebig, an acid is a hydrogen-containing substance in which the hydrogen could be replaced by a metal.[3] Liebig's definition, while completely empirical, remained in use for almost 50 years until the adoption of the Arrhenius definition.[4] [edit] Arrhenius definition 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 for "recognition of the extraordinary services... rendered to the advancement of chemistry by his electrolytic theory of dissociation".[5] As defined at the time of discovery, acid-base reactions are characterized by Arrhenius acids, which dissociate in aqueous solution form hydrogen ions (H+) later recognized to be actually hydronium (H3O+) ions,[5] and Arrhenius bases which form hydroxide (OH−) ions. More recent IUPAC recommendations now suggest the newer term "hydronium"[6] be used in favor of the older accepted term "oxonium"[7] to illustrate reaction mechanisms such as those defined in the Brønsted-Lowry and solvent system definitions more clearly, with the Arrhenius definition serving as a simple general outline of acid-base character.[5] The Arrhenius definition can be summarised as "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 hydrogen ions and hydroxide ions from the dissociation of an acid and base in aqueous solution:
(In modern times, the use of H+ is regarded as a shorthand for H3O+, since it is now known that the bare proton H+ does not exist as a free species in solution.) This 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.[5] In other words, this is a neutralization reaction.
The positive ion from a base forms 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.
[edit] Brønsted-Lowry definitionMain article: Brønsted–Lowry acid-base theory 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—that is, the ability of acids to "donate" hydrogen ions (H+) or protons to bases, which "accept" them.[8] Unlike the Arrhenius definition, the Brønsted-Lowry definition does not refer to the formation of salt and water, but instead to the formation of conjugate acids and conjugate bases, produced by the transfer of a proton from the acid to the base.[5][8] In this definition, an acid is a compound that can donate a proton, and a base is a compound that can receive a proton. An acid-base reaction is, thus, the removal of a hydrogen ion from the acid and its addition to the base.[9] This does not refer to the removal of a proton from the nucleus of an atom, which would require levels of energy not attainable through the simple dissociation of acids, but to removal of a hydrogen ion (H+). The removal of a proton (hydrogen ion) from an acid produces its conjugate base, which is the acid with a hydrogen ion removed, and the reception of a proton by a base produces its conjugate acid, which is the base with a hydrogen ion added. For example, the removal of H+ from hydrochloric acid (HCl) produces the chloride ion (Cl−), the conjugate base of the acid:
The addition of H+ to the hydroxide ion (OH−), a base, produces water (H2O), its conjugate acid:
Thus, the Brønsted-Lowry definition encompasses the Arrhenius definition, but also extends the concept of acid-base reactions to systems in which water is not involved, such as the protonation of ammonia, a base, to form the ammonium ion, its conjugate acid:
This reaction may proceed in the absence of water, such as in the reaction of ammonia with acetic acid:
This definition also provides a theoretical framework for explaining the spontaneous dissociation of water into low concentrations of hydronium and hydroxide ions:
Water, being amphoteric, can act as both an acid and a base; here, one molecule of water acts as an acid, donating a H+ ion and forming the conjugate base, OH−, and a second molecule of water act as a base, accepting the H+ ion and forming the conjugate acid, H3O+.  Thus, the general formula for acid-base reactions according to the Brønsted-Lowry definition is:
where AH represents the acid, B represents the base, and BH+ represents the conjugate acid of B, and A− represents the conjugate base of AH. [edit] Lewis definitionFurther information: Lewis acids and bases The Lewis definition of acid-base reactions, devised by Gilbert N. Lewis in 1923[10] is a further generalization that encompasses the Brønsted-Lowry definition and the solvent-system definitions.[11] Instead of defining acid-base reactions in terms of protons or other bonded substances, the Lewis definition defines a base (referred to as a Lewis base) to be a compound that can donate an electron pair, and an acid (a Lewis acid) to be compound that can receive this electron pair.[11] For example, consider this classical aqueous acid-base reaction:
The Lewis definition does not regard this reaction as the formation of salt and water or the transfer of H+ from HCl to OH−. Instead, it regards the acid to be the H+ ion itself, and the base to be the OH− ion, which has an unshared electron pair. Therefore, the acid-base reaction here, according to the Lewis definition, is the donation of the electron pair from OH− to the H+ ion. This forms a covalent bond between H+ and OH−, thus producing water (H2O). By treating acid-base reactions in terms of electron pairs instead of specific substances, the Lewis definition can be applied to reactions that do not fall under other definitions of acid-base reactions. For example, a silver cation behaves as an acid with respect to ammonia, which behaves as a base, in the following reaction:
The result of this reaction is the formation of an ammonia-silver adduct. In reactions between Lewis acids and bases, there is the formation of an adduct[11] when the highest occupied molecular orbital (HOMO) of a molecule, such as NH3 with available lone electron pair(s) donates lone pairs of electrons to the electron-deficient molecule's lowest unoccupied molecular orbital (LUMO) through a co-ordinate covalent bond; in such a reaction, the HOMO-interacting molecule acts as a base, and the LUMO-interacting molecule acts as an acid.[11] In highly-polar molecules, such as boron trifluoride (BF3),[11] the most electronegative element pulls electrons towards its own orbitals, providing a more positive charge on the less-electronegative element and a difference in its electronic structure due to the axial or equatorial orbiting positions of its electrons, causing repulsive effects from lone pair-bonding pair (Lp-Bp) interactions between bonded atoms in excess of those already provided by bonding pair-bonding pair (Bp-Bp) interactions.[11] Adducts involving metal ions are referred to as co-ordination compounds.[11] [edit] Solvent-system definitionThis definition is based on a generalization of the earlier Arrhenius definition to all auto-dissociating 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, water and ammonia undergo dissociation into hydronium and hydroxide, and ammonium and amide, respectively:
Some aprotic systems also undergo such dissociation, such as dinitrogen tetroxide into nitrosonium and nitrate, and antimony trichloride into dichloroantimonium and 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. The non-aqueous acid-base reactions in liquid ammonia are similar to the reactions in water:
Nitric acid can be a base in liquid sulfuric acid:
The unique strength of this definition shows in describing the reactions in aprotic solvents, for example in liquid N2O4:
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. [edit] Other acid-base theories[edit] Usanovich definitionThe most general definition is that of the Russian chemist Mikhail Usanovich, and can 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. [edit] Lux-Flood definitionThis definition, proposed by German chemist Hermann Lux[12][13] in 1939, further improved by Håkon Flood circa 1947[14] and now commonly used in modern geochemistry and electrochemistry of molten salts, describes an acid as an oxide ion acceptor and a base as an oxide ion donor. For example:[15]
[edit] Pearson definitionMain article: HSAB theory In 1963[16] Ralph Pearson proposed an advanced qualitative concept known as Hard Soft Acid Base principle, later made quantitative with help of Robert Parr in 1984. 'Hard' applies to species that are small, have high charge states, and are weakly polarizable. 'Soft' applies to species that are large, have low charge states and are strongly polarizable. Acids and bases interact, and the most stable interactions are hard-hard and soft-soft. This theory has found use in organic and inorganic chemistry. [edit] See also
[edit] Notes
[edit] References
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