Historic development The concept of an acid–base reaction was first proposed in 1754 by
Guillaume-François Rouelle, who introduced the word "
base" into chemistry to mean a substance which reacts with an acid to give it solid form (as a salt). Bases are mostly bitter in nature.
Lavoisier's oxygen theory of acids The first scientific concept of acids and bases was provided by
Lavoisier in around 1776. Since Lavoisier's knowledge of
strong acids was mainly restricted to
oxoacids, such as (
nitric acid) and (
sulfuric acid), which tend to contain central atoms in high
oxidation states surrounded by oxygen, 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" (). The Lavoisier definition held for over 30 years, until the 1810 article and subsequent lectures by
Sir Humphry Davy in which he proved the lack of oxygen in
hydrogen sulfide (),
hydrogen telluride (), and the
hydrohalic acids. However, Davy failed to develop a new theory, concluding that "acidity does not depend upon any particular elementary substance, but upon peculiar arrangement of various substances". One notable modification of oxygen theory was provided by
Jöns Jacob Berzelius, who stated that acids are oxides of nonmetals while bases are oxides of metals.
Liebig's hydrogen theory of acids In 1838,
Justus von Liebig proposed that an acid is a hydrogen-containing compound whose hydrogen can be replaced by a metal. This redefinition was based on his extensive work on the chemical composition of
organic acids, finishing the doctrinal shift from oxygen-based acids to hydrogen-based acids started by Davy. Liebig's definition, while completely empirical, remained in use for almost 50 years until the adoption of the Arrhenius definition.
Arrhenius definition The first modern definition of acids and bases in molecular terms was devised by
Svante Arrhenius. A hydrogen theory of acids, it followed from his 1884 work with
Friedrich Wilhelm Ostwald in establishing the presence of ions in
aqueous solution and led to Arrhenius receiving the
Nobel Prize in Chemistry in 1903. As defined by Arrhenius: • An
Arrhenius acid is a substance that
ionises in water to form
hydrogen cations (); that is, an acid increases the concentration of H+ ions in an aqueous solution. This causes the
protonation of water, or the creation of the
hydronium () ion. Thus, in modern times, the symbol is interpreted as a shorthand for , because it is now known that a bare proton does not exist as a free species in aqueous solution. This is the species which is measured by
pH indicators to measure the acidity or basicity of a solution. • An
Arrhenius base is a substance that dissociates in water to form
hydroxide () ions; that is, a base increases the concentration of ions in an aqueous solution. The Arrhenius definitions of
acidity and
alkalinity are restricted to aqueous solutions and are not valid for most non-aqueous solutions, and refer to the concentration of the solvent ions. Under this definition, pure and HCl dissolved in toluene are not acidic, and molten NaOH and solutions of calcium amide in liquid ammonia are not alkaline. This led to the development of the Brønsted–Lowry theory and subsequent Lewis theory to account for these
non-aqueous exceptions. The reaction of an acid with a base is called a
neutralization reaction. The products of this reaction are a
salt and water. \text{acid} \ + \ \text{base} \ \longrightarrow \ \text{salt} \ + \ \text{water} In this traditional representation an acid–base neutralization reaction is formulated as a
double-replacement reaction. For example, the reaction of
hydrochloric acid (HCl) with
sodium hydroxide (NaOH) solutions produces a solution of
sodium chloride (NaCl) and some additional water molecules. \ce{HCl_{(aq)} {} + NaOH_{(aq)} -> NaCl_{(aq)} {} + H2O} The modifier (
aq) in this equation was implied by Arrhenius, rather than included explicitly. It indicates that the substances are dissolved in water. Though all three substances, HCl, NaOH and NaCl are capable of existing as pure compounds, in
aqueous solutions they are fully dissociated into the aquated ions and .
Example: Baking powder Baking powder is used to cause the dough for breads and cakes to "rise" by creating millions of tiny
carbon dioxide bubbles. Baking powder is not to be confused with
baking soda, which is
sodium bicarbonate (). Baking powder is a mixture of baking soda (sodium bicarbonate) and acidic salts. The bubbles are created because, when the baking powder is combined with water, the sodium bicarbonate and acid salts react to produce gaseous
carbon dioxide. Whether commercially or domestically prepared, the principles behind baking powder formulations remain the same. The acid–base reaction can be generically represented as shown: \ce{NaHCO3 + H+ -> Na+ + CO2 + H2O} The real reactions are more complicated because the acids are complicated. For example, starting with sodium bicarbonate and
monocalcium phosphate (), the reaction produces carbon dioxide by the following
stoichiometry: \ce{14 NaHCO3 + 5 Ca(H2PO4)2 -> 14 CO2 + Ca5(PO4)3OH + 7 Na2HPO4 + 13 H2O} ("MCP") is a common acid component in domestic baking powders. A typical formulation (by weight) could call for 30% sodium bicarbonate, 5–12%
monocalcium phosphate, and 21–26%
sodium aluminium sulfate. Alternately, a commercial baking powder might use
sodium acid pyrophosphate as one of the two acidic components instead of sodium aluminium sulfate. Another typical acid in such formulations is
cream of tartar (), a derivative of
tartaric acid. is based upon the idea of
protonation of bases through the
deprotonation of acids – that is, the ability of acids to "donate"
hydrogen cations () otherwise known as
protons to bases, which "accept" them. An acid–base reaction is, thus, the removal of a proton from the acid and its addition to the base. The removal of a proton from an acid produces its
conjugate base, which is the acid with a proton removed. The reception of a proton by a base produces its
conjugate acid, which is the base with a proton added. Unlike the previous definitions, the Brønsted–Lowry definition does not refer to the formation of salt and solvent, but instead to the formation of
conjugate acids and
conjugate bases, produced by the transfer of a proton from the acid to the base. in the same year as Brønsted–Lowry, but it was not elaborated by him until 1938. For example,
boron trifluoride, is a typical Lewis acid. It can accept a pair of electrons as it has a vacancy in its
octet. The
fluoride ion has a full octet and can donate a pair of electrons. Thus \ce{BF3 + F- -> BF4-} is a typical Lewis acid, Lewis base reaction. All compounds of
group 13 elements with a formula can behave as Lewis acids. Similarly, compounds of
group 15 elements with a formula , such as
amines, , and
phosphines, , can behave as Lewis bases.
Adducts between them have the formula with a
dative covalent bond, shown symbolically as ←, between the atoms A (acceptor) and D (donor). Compounds of
group 16 with a formula may also act as Lewis bases; in this way, a compound like an
ether, , or a
thioether, , can act as a Lewis base. The Lewis definition is not limited to these examples. For instance,
carbon monoxide acts as a Lewis base when it forms an adduct with boron trifluoride, of formula . Adducts involving metal ions are referred to as co-ordination compounds; each ligand donates a pair of electrons to the metal ion. Germann pointed out that in many solutions, there are ions in equilibrium with the neutral solvent molecules: •
solvonium ions: a generic name for positive ions. These are also sometimes called solvo-acids; when
protonated solvent, they are
lyonium ions. •
solvate ions: a generic name for negative ions. These are also sometimes called solve-bases; in 1939, further improved by
Håkon Flood and is still used in modern
geochemistry and
electrochemistry of
molten salts. This definition describes an acid as an oxide ion () acceptor and a base as an oxide ion donor. For example: \begin{array}{ccccl} _\text{(base)} & & _\text{(acid)} \\[4pt] \ce{MgO} &+& \ce{CO2} &\longrightarrow& \ce{MgCO3} \\[4pt] \ce{CaO} &+& \ce{SiO2} &\longrightarrow& \ce{CaSiO3} \\[4pt] \ce{NO3-} &+& \ce{S2O7^2-} \!\! &\longrightarrow& \ce{NO2+ + 2 SO4^2-} \end{array} This theory is also useful in the systematisation of the reactions of
noble gas compounds, especially the xenon oxides, fluorides, and oxofluorides.
Usanovich definition Mikhail Usanovich developed a general theory that does not restrict acidity to hydrogen-containing compounds, but his approach, published in 1938, was even more general than Lewis theory. Usanovich's theory can be summarized as defining an acid as anything that accepts negative species or donates positive ones, and a base as the reverse. This defined the concept of
redox (oxidation-reduction) as a special case of acid–base reactions. Some examples of Usanovich acid–base reactions include: \begin{array}{ccccll} _\text{(base)} & & _\text{(acid)} \\[4pt] \ce{Na2O} &+& \ce{SO3} &\longrightarrow& \ce{2Na+ {} + \ SO4^2-} & \text{(species exchanged: } \ce{O^2-} \text{anion)} \\[4pt] \ce{3(NH4)2S} &+& \ce{Sb2S5} &\longrightarrow& \ce{6 NH4+ {} + \ 2 SbS4^3-} & \text{(species exchanged: } \ce{3 S^2-} \text{ anions)} \\[4pt] \ce{2Na} &+& \ce{Cl2} &\longrightarrow& \ce{2 Na+ {} + \ 2 Cl-} & \text{(species exchanged: 2 electrons)} \end{array} ==Rationalizing the strength of Lewis acid–base interactions==