A brief review of the various concepts regarding acids and bases are given here.
THE ARRHENIUS (CLASSICAL) CONCEPT (1884)
According to this an acid is defined as a hydrogen containing substance which gives 114 ions (i.e., H30+ hydronium ions) when dissolved in water.
A base is a substance which coliains OH groups and gives hydroxyl ions OH – when dissolved in water.
Arrhenius concept is based upon ionic dissociation of compound in water. For example, HCl is an acid because it produces H30+ ions in water but CH4 is not. Similarly, NaOH is a base because it furnishes OH- ions, whereas C2H5OH is not a base.
HC1 + H2O > H30+ + Cl-
NaOH + H2O > Na+ + OH
The process of neutralization of an acid by a base can be represented by the reaction to form neutral water.
H+ + OH > H2O
With this concept, many aspects of acid-base behavior were understood. For instance, the constant heat of neutralization of a strong acid by a strong base can readily be explained in terms of Arrhenius concept because the reaction ‘involves only the combination of a hydrogen ion and a hydroxyl ion in all such neutralization reactions.
It explains the catalytic properties of acids. Arrhenius theory affords a correlation between the electrolytic dissociation and the concentrations of the hydrogen ion. The mobility of the hydrogen ions parallel the catalytic activity of the solution if the hydrogen ion is truly the source of the catalytic properties.
According to this theory, all the acid-base reactions are limited to aqueous medium only. It does not explain the acid-base reactions taking place in non-aqueous solvents such as liquid ammonia.
It also cannot explain the reactions in gas phase where no solvent is present.
Similarly, the definition of a base under this concept is restricted to compounds containing hydroxyl ions only, whereas many organic compounds as well as ammonia which exhibit basic properties cannot be explained by this definition. Similarly, there are many acidic compounds which do not contain hydrogen and cannot be explained on the basis of Arrhenius concept. Hence new concepts were put forward to explain more general cases of acids and bases.
THE PROTONIC OR LOWRY-BRONSTED CONCEPT (1923)
According to Bronsted:
An acid is defined as a species (a compound or an ion) which donates or tends to donate a proton (II+ ion).
A base is a species which accepts or tends to accept a proton. Acid-base reaction is the transfer of a proton from an acid to a base. The dissociation of an acid HA can be represented as:
HA > A- + H+
Acid Base proton CH3COOH > CH3C00- +h
Acid Base Proton
According to this definition, any negatively charged ion (anion) acts as a base. Thus, CH3C00- is a base and is said to be conjugate base of acetic acid In an acid-base reaction, an acid yields a base (conjugate) and base after accepting proton yields a conjugate acid.
The acid-base reaction is represented as:
Al Ba > B1 A2
Bronsted Conjugate acid base & base acid
The conjugate acid-base pairs are species on opposite sides of an equation that differ by a proton. The weaker acids have stronger conjugate base pairs and stronger acids have weaker conjugate base.
Thus Cl S024 , OH – are conjugate bases of HC1, HSO4 and H2O, respectively. Similarly, H2O, HSO4 and HC1 are conjugate acids of the bases OW, S024 and Cl -, respectively.
The following species may be regarded as acids:
Molecular Species: HC1, H2SO4, CH3COOH, HCN, H2S, H2O etc.
Anionic Species:HSO4, HCO3, H2PO4, HP024, HS etc.
Cationic Species: H30+, NHS, [Cu(H20)4]2+ etc.
The following species may be regarded as bases:
Molecular Species: H2O, NH3, CH3NH2 etc.
Anionic Species: OH-, HS -, S2-, HCO3, HSO4, C1- etc.
Cationic Species: [Fe(H20)50H]2+, [Cu(H20)3 Of1]+ etc.
From the above examples, it is found that some of the species act both as acids and bases depending upon the manner they behave in the given reaction. Amphiprotic Species: A species that acts bath as a proton donor and a proton acceptor is said to be amphiprotic. For example:
(i) H2O is amphiprotic. It loses proton to a base such as NH3 or accepts a proton from an acid such as HC1.
H20 + NH3 > NH+4 + OH-
H20 + HCl > H30+ + Cl-
The proton-containing negative ions are amphiprotic. For example:
HS- + OH – > S-2 + H2O
HS- + H30+ > H2S + H2O
The conjugate base and acid are shown as:
OH- Conjugate base of water
— H+ H2O +H+
H30+ Conjugate acid of water CO3 Conjugate base of HCO3
H2CO3 Conjugate acid of HCO3
Amphoteric hydroxides react with both acids and bases becadse they are equally amphiprotic.
Al(H20)3 (OH)3 + OH – > Al(H20)2 (OH)4 + H2O
Al(H20)3 (OH)3 + H2O+ > Al(H20)4 (OH)2 + H2O
Polyprotic Acids: Acids containing one proton which can be donated are monoprotic acids. Those acids which contain more than one donatable proton are known as “Polyprotic acids” e.g., HC1, HNO3, HCN etc. are monoprotic, whereas H3PO4, H2SO4, H3AsO4 are Polyprotic.
Non-aqueous solutions also lose or gain protons and fit into the Bronsted acids and bases. For example, water is dissolved in liquid ammonia.
H20 + NH3 > OH- + NH+4