Acid–base definitions
The concept of an acid-base reaction was number one proposed in 1754 by Guillaume-François Rouelle, who gave the word "base" into chemistry to mean a substance which reacts with an acid to dispense it solid make-up as a salt. Bases are mostly bitter in nature.
The first scientific concept of acids and bases was provided by Lavoisier in around 1776. Since Lavoisier's cognition of strong acids was mainly restricted to oxoacids, such(a) as nitric acid and SO sulfuric acid, which tend to contain central atoms in high oxidation states surrounded by oxygen, and since he was non 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 ὀξύς oxys meaning "acid" or "sharp" and γεινομαι geinomai meaning "engender". 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 S, H2Te, and the hydrohalic acids. However, Davy failed to develop a new theory, concluding that "acidity does not depend upon all particular elementary substance, but upon peculiar arrangement of various substances". One notable correct of oxygen theory was provided by Jöns Jacob Berzelius, who stated that acids are oxides of nonmetals while bases are oxides of metals.
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 develope 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 totally empirical, remained in usage for near 50 years until the adoption of the Arrhenius definition.
The first contemporary 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:
This causes the protonation of water, or the creation of the hydronium H3O+ ion. Thus, in modern times, the symbol H+ is interpreted as a shorthand for H3O+, because it is now required that a bare proton does not make up as a free category in aqueous solution. it is species which is measured by pH indicators to measure the acidity or basicity of a 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 H2SO4 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 developing 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.
In this traditional explanation 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 written of sodium chloride, NaCl, and some extra water molecules.
The modifier aq in this equation was implied by Arrhenius, rather than target 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 H+, Cl−, Na+ and OH−.
The Brønsted–Lowry definition, formulated in 1923, independently by Johannes Nicolaus Brønsted in Denmark and Martin Lowry in England, is based upon the idea of protonation of bases through the deprotonation of acids – that is, the ability of acids to "donate" hydrogen ions H+—otherwise known as protons—to bases, which "accept" them.
An acid–base reaction is, thus, the removal of a hydrogen ion from the acid and its addition to the base. The removal of a hydrogen ion from an acid produces its conjugate base, which is the acid with a hydrogen ion removed. The reception of a proton by a base produces its conjugate acid, which is the base with a hydrogen ion added.
Unlike the preceding definitions, the Brønsted–Lowry definition does not refer to the configuration of salt and solvent, but instead to the layout of conjugate acids and conjugate bases, produced by the transfer of a proton from the acid to the base. In this approach, acids and bases are fundamentally different in behavior from salts, which are seen as electrolytes, noted to the theories of Debye, Onsager, and others. An acid and a base react not to produce a salt and a solvent, but to form a new acid and a new base. The concept of neutralization is thus absent. Brønsted–Lowry acid–base behavior is formally self-employed adult of any solvent, creating it more all-encompassing than the Arrhenius model. The solution of pH under the Arrhenius good example depended on alkalis bases dissolving in water aqueous solution. The Brønsted–Lowry model expanded what could be pH tested using insoluble and soluble solutions gas, liquid, solid.
The general formula for acid–base reactions according to the Brønsted–Lowry definition is:
where HA represents the acid, B represents the base, BH+ represents the conjugate acid of B, and A− represents the conjugate base of HA.
For example, a Brønsted-Lowry model for the dissociation of hydrochloric acid HCl in aqueous solution would be the following:
The removal of H+ from the HCl produces the chloride ion, Cl−, the conjugate base of the acid. The addition of H+ to the H2O acting as a base forms the hydronium ion, H3O+, the conjugate acid of the base.
Water is amphoteric—that is, it can act as both an acid and a base. The Brønsted-Lowry model explains this, showing the dissociation of water into low concentrations of hydronium and hydroxide ions:
This equation is demonstrated in the image below:
Here, one molecule of water acts as an acid, donating an H+ and forming the conjugate base, OH−, and amolecule of water acts as a base, accepting the H+ ion and forming the conjugate acid, H3O+.
As an example of water acting as an acid, consider an aqueous solution of pyridine, C5H5N.
In this example, a water molecule is split into a hydrogen ion, which is donated to a pyridine molecule, and a hydroxide ion.
In the Brønsted-Lowry model, the solvent does not necessarily have to be water, as is required by the Arrhenius Acid-Base model. For example, consider what happens when acetic acid, CH3COOH, dissolves in liquid ammonia.
An H+ ion is removed from acetic acid, forming its conjugate base, the acetate ion, CH3COO−. The addition of an H+ ion to an ammonia molecule of the solvent creates its conjugate acid, the ammonium ion, .
The Brønsted–Lowry model calls hydrogen-containing substances like HCl acids. Thus, some substances, which many chemists considered to be acids, such as SO3 or BCl3, are excluded from this kind due to lack of hydrogen. Gilbert N. Lewis wrote in 1938, "To restrict the multinational of acids to those substances that contain hydrogen interferes as seriously with the systematic apprehension of chemistry as would the restriction of the term oxidizing agent to substances containing oxygen." Furthermore, KOH and KNH2 are not considered Brønsted bases, but rather salts containing the bases OH− and .
The hydrogen requirement of Arrhenius and Brønsted–Lowry was removed by the Lewis definition of acid–base reactions, devised by Gilbert N. Lewis in 1923, in the same year as Brønsted–Lowry, but it was not elaborated by him until 1938. 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 a compound that can get this electron pair.
For example, boron trifluoride, BF3 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
is a typical Lewis acid, Lewis base reaction. All compounds of group 13 elements with a formula AX3 can behave as Lewis acids. Similarly, compounds of group 15 elements with a formula DY3, such as amines, NR3, and phosphines, PR3, can behave as Lewis bases. Adducts between them have the formula X3A←DY3 with a dative covalent bond, shown symbolically as ←, between the atoms A acceptor and D donor. Compounds of group 16 with a formula DX2 may also act as Lewis bases; in this way, a compound like an ether, R2O, or a thioether, R2S, 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 F3B←CO.
Adducts involving metal ions are referred to as co-ordination compounds; used to refer to every one of two or more people or things ligand donates a pair of electrons to the metal ion. The reaction
can be seen as an acid–base reaction in which a stronger base ammonia replaces a weaker one water
The Lewis and Brønsted–Lowry definitions are consistent with each other since the reaction
is an acid–base reaction in both theories.
One of the limitations of the Arrhenius definition is its reliance on water solutions. Edward Curtis Franklin studied the acid–base reactions in liquid ammonia in 1905 and pointed out the similarities to the water-based Arrhenius theory. Albert F.O. Germann, working with liquid phosgene, , formulated the solvent-based theory in 1925, thereby generalizing the Arrhenius definition to continue aprotic solvents.
Germann pointed out that in numerous solutions, there are ions in equilibrium with the neutral solvent molecules:
For example, water and amide, respectively:
Some aprotic systems also undergo such dissociation, such as dinitrogen tetroxide into nitrosonium and nitrate, antimony trichloride into dichloroantimonium and tetrachloroantimonate, and phosgene into chlorocarboxonium and chloride:
A solute that causes an put in the concentration of the solvonium ions and a decrease in the concentration of solvate ions is defined as an acid. A solute that causes an increase in the concentration of the solvate ions and a decrease in the concentration of the solvonium ions is defined as a base.
Thus, in liquid ammonia, supplying is a strong base, and NO supplying is a strong acid. In liquid sulfur dioxide , thionyl compounds supplying behave as acids, and sulfites supplying 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 O:
Because the solvent system definition depends on the solute as alive as on the solvent itself, a particular solute can be either an acid or a base depending on the pick of the solvent: is a strong acid in water, a weak acid in acetic acid, and a weak base in fluorosulfonic acid; this characteristic of the theory has been seen as both a strength and a weakness, because some substances such as and have been seen to be acidic or basic on their own right. On the other hand, solvent system theory has been criticized as being too general to be useful. Also, it has been thought that there is something intrinsically acidic about hydrogen compounds, a property not divided up by non-hydrogenic solvonium salts.
This acid–base theory was a revival of the oxygen theory of acids and bases proposed by German chemist Hermann Lux in 1939, further renovation by Håkon Flood circa 1947 and is still used in sophisticated 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:
This theory is also useful in the systematisation of the reactions of noble gas compounds, especially the xenon oxides, fluorides, and oxofluorides.
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 effect of acid–base reactions
Some examples of Usanovich acid–base reactions include: