Acids And Bases: Theories
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The substances are classified into electrolytes and non-electrolytes based on their ability to conduct electricity. The substances like sodium chloride which conduct electricity in their aqueous solution are called electrolytes. The molecular compounds like sugar which do not conduct electricity are known as non-electrolytes. Electrolytes further into strong and weak electrolytes. Strong electrolytes: Strong electrolytes undergo dissociation completely when dissolved in water. EX: HCl is considered a strong electrolyte as it undergoes dissociation completely when dissolved in water. HCl(aq) → H+(aq) + Cl-(aq) Weak electrolytes: Weak electrolytes undergo dissociation only to a limited extent when dissolved in water. EX: acetic acid is considered a weak electrolyte as it undergoes ionisation to a very limited extent. CH3COOH(aq)  ⇌ H+(aq) + CH3COO-(aq) Ionic equilibrium: The equilibrium involving the ions in aqueous solution of weak electrolytes is called ionic equilibrium. Ionic equilibrium can be conveniently studied under two headings: Dissociation equilibria Solubility equilibria Dissociation equilibria: The equilibrium between a dissolved undissociated molecule and its ions is known as dissociation equilibria. In general, dissociation equilibria involve either acids or bases. Acids: They are sour to the taste. They turn blue litmus red. They react with active metals like zinc, and liberate hydrogen gas. Bases: They are bitter in taste. They turn red litmus blue. They are soapy to touch. Theories of Acids and Bases: Arrhenius theory: According to his theory an acid is a substance that contains hydrogen and ionises in an aqueous solution to give hydrogen ions. EX: HCl contains hydrogen and ionises in water to H+ and Cl- ions. HCl(aq) → H+(aq) + Cl-(aq) H+ is very reactive and cannot exist independently in aqueous solutions. Instead, it bonds to an oxygen atom of the water molecule and forms H3O plus, the hydronium ion. HCl(aq) → H3O+(aq) + Cl-(aq) A base is a substance that contains the hydroxyl group and ionises in an aqueous solution to give hydroxide ions. Ex: NaOH contains a hydroxyl group and ionises in water to Na plus and OH minus ions. NaOH(aq) → Na+(aq) + OH-(aq) Limitations: It is limited to aqueous solutions only. Some substances exhibit the properties of a base though they do not contain hydroxide ions. EX: Liquid ammonia. Bronsted-Lowry theory: According to this theory, 'an acid is a substance that exhibits a tendency to donate a proton and a base is a substance that exhibits a tendency to gain a proton'. Proton donors are acids and proton acceptors are bases. Consider a general equation HA(aq) + H2O(l) ⇌ H3O+(aq) + A-(aq) In the forward reaction, HA donates proton to water, and thus, acts as a Bronsted-Lowry acid, while water accepts the proton, and acts as a Bronsted-Lowry base. A related pair of an acid and a base which differ by a single proton is called a conjugate acid-base pair. HA       +          H2O        ⇌          H3O+        +             A- acid               base                    conjugate acid          conjugate base Limitations: The theory fails to explain some reactions that do not involve a proton transfer. It fails to explain the acidic behaviour of electron-deficient molecules like AlCl3 and BCl3. Lewis theory: According to this theory, an acid is any molecule or ion that can accept an electron pair to form a coordinate covalent bond with the donor. EX: H+, BF3, AlCl3 and SnCl4 A base is any molecule or ion that can donate a pair of electrons to form a coordinate covalent bond with the acceptor. Ex: NH3, H2O and OH- Limitations: This theory fails to explain the strength of acids and bases. All the acid-base reactions do not involve coordinate covalent bond formation Generally, acid-base reactions are very fast, but in certain cases, the formation of a coordinate covalent bond is very slow.

#### Summary

The substances are classified into electrolytes and non-electrolytes based on their ability to conduct electricity. The substances like sodium chloride which conduct electricity in their aqueous solution are called electrolytes. The molecular compounds like sugar which do not conduct electricity are known as non-electrolytes. Electrolytes further into strong and weak electrolytes. Strong electrolytes: Strong electrolytes undergo dissociation completely when dissolved in water. EX: HCl is considered a strong electrolyte as it undergoes dissociation completely when dissolved in water. HCl(aq) → H+(aq) + Cl-(aq) Weak electrolytes: Weak electrolytes undergo dissociation only to a limited extent when dissolved in water. EX: acetic acid is considered a weak electrolyte as it undergoes ionisation to a very limited extent. CH3COOH(aq)  ⇌ H+(aq) + CH3COO-(aq) Ionic equilibrium: The equilibrium involving the ions in aqueous solution of weak electrolytes is called ionic equilibrium. Ionic equilibrium can be conveniently studied under two headings: Dissociation equilibria Solubility equilibria Dissociation equilibria: The equilibrium between a dissolved undissociated molecule and its ions is known as dissociation equilibria. In general, dissociation equilibria involve either acids or bases. Acids: They are sour to the taste. They turn blue litmus red. They react with active metals like zinc, and liberate hydrogen gas. Bases: They are bitter in taste. They turn red litmus blue. They are soapy to touch. Theories of Acids and Bases: Arrhenius theory: According to his theory an acid is a substance that contains hydrogen and ionises in an aqueous solution to give hydrogen ions. EX: HCl contains hydrogen and ionises in water to H+ and Cl- ions. HCl(aq) → H+(aq) + Cl-(aq) H+ is very reactive and cannot exist independently in aqueous solutions. Instead, it bonds to an oxygen atom of the water molecule and forms H3O plus, the hydronium ion. HCl(aq) → H3O+(aq) + Cl-(aq) A base is a substance that contains the hydroxyl group and ionises in an aqueous solution to give hydroxide ions. Ex: NaOH contains a hydroxyl group and ionises in water to Na plus and OH minus ions. NaOH(aq) → Na+(aq) + OH-(aq) Limitations: It is limited to aqueous solutions only. Some substances exhibit the properties of a base though they do not contain hydroxide ions. EX: Liquid ammonia. Bronsted-Lowry theory: According to this theory, 'an acid is a substance that exhibits a tendency to donate a proton and a base is a substance that exhibits a tendency to gain a proton'. Proton donors are acids and proton acceptors are bases. Consider a general equation HA(aq) + H2O(l) ⇌ H3O+(aq) + A-(aq) In the forward reaction, HA donates proton to water, and thus, acts as a Bronsted-Lowry acid, while water accepts the proton, and acts as a Bronsted-Lowry base. A related pair of an acid and a base which differ by a single proton is called a conjugate acid-base pair. HA       +          H2O        ⇌          H3O+        +             A- acid               base                    conjugate acid          conjugate base Limitations: The theory fails to explain some reactions that do not involve a proton transfer. It fails to explain the acidic behaviour of electron-deficient molecules like AlCl3 and BCl3. Lewis theory: According to this theory, an acid is any molecule or ion that can accept an electron pair to form a coordinate covalent bond with the donor. EX: H+, BF3, AlCl3 and SnCl4 A base is any molecule or ion that can donate a pair of electrons to form a coordinate covalent bond with the acceptor. Ex: NH3, H2O and OH- Limitations: This theory fails to explain the strength of acids and bases. All the acid-base reactions do not involve coordinate covalent bond formation Generally, acid-base reactions are very fast, but in certain cases, the formation of a coordinate covalent bond is very slow.

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