Arrhenius Theory of Acids and Bases
It was proposed by Arrhenius to explain
the behaviour of acids and bases in aqueous solutions.
Arrhenius acids:— A substance that produces hydrogen
ions (H⁺) when dissolved in water.
ex:
(i) HCl (aq) → H⁺ (aq) + Cl⁻ (aq)
(ii) H₂SO₄ (aq) → 2H⁺ (aq) + SO₄²⁻ (aq)
(iii) HNO₃ (aq) → H⁺ (aq) + NO₃⁻ (aq)
Arrhenius base:— A substance that produces
hydroxide ions (OH⁻) when dissolved in water.
ex:
(i) NaOH (aq) → Na⁺ (aq) + OH⁻ (aq)
(ii) KOH (aq) → K⁺ (aq) + OH⁻ (aq)
(iii) Ca(OH)₂ (aq) → Ca²⁺ (aq) + 2OH⁻ (aq)
Limitations:—
i. This theory is applicable only to aqueous solutions
but not to non–aqueous solutions.
ii. It could not explain the acidic nature of CO₂,
SO₃, Cl₂O, N₂O₅ and basic nature of NH₃,
CaO, Na₂O in non aqueous medium but it
can explain in aqueous medium.
iii. This theory could not explain acidic nature of
BF₃, BCl₃ and AlCl₃.
Bronsted–Lowry Theory:—
It was proposed by Brønsted–Lowry theory to explain
acid–base behaviour based on proton (H⁺) transfer.
This theory is applicable to reactions occurring even
in non–aqueous media.
Bronsted–Lowry acid:— The substance which donates
a proton to another substance is called acid. (donor proton)
ex: HCl, HNO₃, H₂SO₄, CH₃COOH, H₃PO₄ …
Bronsted–Lowry base:— The substance which can
accept proton is called a base. (accept proton)
acceptor is a base.
ex: NH₃, OH⁻, HSO₄⁻, Cl⁻ …
An acid can donate a proton only if a base
is present to accept it. Every acid has a conjugate
base, formed after loss of proton.
ex: acid — conjugate base
-
HCl — Cl⁻
-
H₂SO₄ — HSO₄⁻
-
HNO₃ — NO₃⁻
-
CH₃COOH — CH₃COO⁻
Every base has a corresponding conjugate
acid, formed after gain of a proton.
ex:
Base — Conjugate acid
NH₃ — NH₄⁺
PO₄³⁻ — HPO₄²⁻
OH⁻ — H₂O
H₂O — H₃O⁺ Conjugate acid–base pairs:—
A conjugate acid–base pair differ by
one proton (H⁺).
ex: NH₃ + H₂O ⇌ NH₄⁺ + OH⁻
base acid conjugate acid conjugate base
The conjugate base of strong acid is
weak base and the conjugate base of weak acid is strong base.
ex: HCl is a strong acid and Cl⁻ is a weak base.
CH₃COOH is a weak acid and CH₃COO⁻ is strong base.
The conjugate acid of strong base is weak acid and
the conjugate acid of weak base is strong.
ex: OH⁻ is a strong base and H₂O is a weak acid.
C₂H₅OH is a weak base and (C₂H₅OH₂⁺ )is a strong acid.
Strong acid: An acid which shows greater tendency to donate
protons is a strong acid.
ex: HClO₄, H₂SO₄, HNO₃.
Weak acid: An acid which shows less tendency to donate
protons is a weak acid.
ex: CH₃COOH, HCN, H₃PO₄.
Strong base: A base which shows greater tendency to accept
protons is a strong base.
ex: CH₃COO⁻, CN⁻, H₂PO₄⁻.
Weak base: A base which exhibits less tendency
to accept proton is a weak base.
ex: NO₃⁻, Cl⁻, HSO₄⁻, ClO₄⁻.
Advantages:—
1. It explains acid–base reactions in aqueous and non–aqueous media.
2. It includes substances like NH₃ as bases which are
not explained by Arrhenius theory.
3. It introduces the important concept of conjugate
acid–base pairs.
4. It explains acid–base behaviour in gaseous
phase reactions.
Limitations:—
1. The theory applies only to proton–transfer reactions.
2. It cannot explain acid–base reactions in which no proton
is involved.
3. It could not explain neutralization reactions without
proton transfer reactions.
4. It fails to explain the acidic nature of compounds
like BF₃, AlCl₃ which does not contain hydrogen.
5. It does not account for acid–base reactions involving electron pairs transfer.
LEWIS – ACID BASE THEORY :—
It was proposed by G.N. Lewis in 1923. It provides a broad
and modern concept of acids and bases based on the behaviour
of electrons. According to this theory, acid–base reactions
involve the transfer of an electron pair rather than hydrogen
ions. This theory is useful in understanding coordination
compounds and complex formation.
Lewis acid:— The substance which accepts an electron
pair during a chemical reaction is called lewis acid.
Lewis base:— The substance which donates an electron
pair during a chemical reaction is called lewis base.
Types of Lewis acids:—
1. All simple cations which can able to accept electron
pairs.
ex: H⁺, Fe²⁺, Fe³⁺, Cu²⁺, Ni²⁺, Ag⁺, Co³⁺, Al³⁺ …
2. Atoms having six electrons in their valence shell
ex: O, S, Se, Te.
3. electron deficient molecules.
ex: BCl₃, AlCl₃, BF₃, FeCl₃.
4. molecules having expanded in which central atom
can expand its octet configuration
ex: SF₄, FeCl₃, SnCl₄ …
5. molecules having multiple bonds between
dissimilar atoms.
ex: CO₂, SO₂, SO₃, NO₂ …
Types of Lewis bases:—
1. All anions which can able to donate electron pairs
ex: F⁻, Cl⁻, OH⁻, CN⁻, SCN⁻, NH₂⁻ …
2. molecules having lone pair of electrons which can able
to donate
ex: NH₃, H₂O, R–OH, R₂S, R–NH₂, R–O–R …
3. molecules having multiple bonds between
similar atoms.
ex: H₂C = CH₂, HC ≡ CH.
4. molecules which can act as ligands.
ex: CO, NO.
Limitations:—
(i). It does not clearly explain the relative strength
of acids and bases.
(ii). Acid–base neutralization reactions occur quickly.
But, according to this theory, neutralization reactions
occur slowly.
(iii). It could not explains the catalytic activity
of H⁺ ions.
(iv). It does not distinguish between acid–base
reactions and redox reactions.