CBSE Class 11 Chemistry Notes For Acids And Bases

Acids And Bases

Acids and bases according to Arrhenius’s theory

Acids:’

Hydrogen-containing compounds that ionize in an aqueous solution to produce H+ ions are called acids.

Example:

The hydrogen-containing compounds such as HCl, HNO₃, H₂SO₄, CH₃COOH, etc., ionize in aqueous solutions to form H⁺ ions. Thus, these compounds are acids according to Arrhenius’s theory.

Bases A compound that ionizes in an aqueous solution to produce hydroxyl ions (OH-) is called a base.

Example: The compounds such as NaOH, KOH, Ca(OH)₂ NH₄OH, etc., ionize in aqueous solution to produce OH’ ions and hence are termed as bases according to Arrhenius theory.

1. NaOH(ag) Na⁺(ag) + OH^–(aq)
2. KOH(ag)→ K⁺(aq) + OH^–(aq)
3. Ca(OH)2(aq)→ Ca²⁺(aq) + 2OH^–(aq)
4. NH4OH(aq)→ NH⁺(aq) + OH^–(aq)

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Limitations of Arrhenius theory: Arrhenius theory is useful for defining acids and bases. It explains the acid-base neutralization reaction by the simple equation:

H₃O+(aq) + OH-(aq)→2H₂O(Z). However, there are certain limitations of this theory.

According to this theory, the presence of water is essential for a compound to exhibit its acidic or basic properties. However, the fact that acidic or basic property of a compound is its characteristic property, which is independent of the presence of water.

• The acidic or basic properties of a substance that is insoluble in water cannot be explained by this theory.
• The acidity or basicity of any compound in non-aqueous solvents cannot be explained by Arrhenius’s theory. For example, the acidity of NH₄C1 orbasicity of NaNH₂ in liquid
  ammonia cannot be explained by this theory.
• According to Arrhenius’s theory, compounds containing only hydroxyl ions are considered as bases. Consequently, the basicity of ammonia (NH₃), methylamine (CH₃NH₂), aniline (C₆H₅NH₂), etc., cannot be explained by this theory.
• According to Arrhenius’s theory, only the hydrogen-containing compounds that ionize in aqueous solution to produce H+ ions are considered acids. Consequently, the acidity of compounds such as PC1₅, BF₃, and A1C1₃ cannot be explained by this theory.

Acids And Bases According To Bronsted Lowry Concept (Protonic Theory)

Definitions of acids and- bases according to the theory proposed by J.N. Bronsted and T.M. Lowry are given below:

1. Acid: An acid is a substance that can donate a proton (or H+ ion)
2. Base: A base is a substance that can accept a proton (or J+ ions)

So, according to this theory, an acid is a proton donor and a base is a proton acceptor.

Example: HCl(aq) + H₂O(aq)→H₃O⁺(aq) + Cl^–(aq)

In this reaction, HC1 donates one proton, behaving as acid, while H20 accepts a proton, behaving as a base.

According to this theory, apart from the mill compounds (HCI, HNO₃, CH₃COOH, etc cations example;

NH⁴⁺, C₆H₅NH₃⁺,[Fe(H₂O)₆]³⁺, [A](H₂O)₆]3⁺, etc.) and anions (example HSO^–₄, HCO^–₃, HC₂O^–₄ etc.] and The acidic properties of these three types of substances are shown by the following reactions ;

CH₃COOOH(aq)+ H₂O(l) ⇌ H₃O⁺(aq)+CH₃COO^–(aq)

H₂SO₄(aq) + 2H₂O(l)⇌ 2H₃O⁺(aq)+SO₄^2-(aq)

NH₄⁺(aq) + H₂O(l)⇌ H₃O⁺(aq) +NH₃(aq)

[Fe(H₂O)₆]³⁺ (aq) + H₂O(l) ⇌ [Fe(H₂O)₅OH]²⁺(aq)+ H₃O⁺(aq)

H₂PO₄^–(aq) + H₂O(l) ⇌ H₃O⁺(aq)+HPO₄^2-(aq)

Similarly, apart from the neutral compounds (e.g., NH₃, C₆H5NH₂, H₂O, etc.), a large number of unions

Example: OH^–, CH₃COO^–, CO₃^2-, etc.) can act as a base,

The following reactions indicate the basic properties of these two lands of substances:

1. CH₃COO^–(aq) + H₂O(l)⇌ CH₃COOH(aq) + OH^–(aq)
2. NH₃(aq) +H₂O(l)⇌NH⁺4(aq) + OH^–(aq)
3. HPO₄^2--(aq) + H₂O(l)⇌ (aq) + OH^–(aq)

Concept of conjugate acid-base pair:

Concept of conjugate acid-base pair Definition:

Apair of species (a neutral compound and the Ion produced from it or, an ion and a neutral compound formed from it or, an ion and the other ion produced from it) having a difference of one proton is called a conjugate acid-base pair.

Examples: (H₂O, H₃O⁺), (H₂PO₄^–, HPO₄^2-), (CH₃COO^–, CH₃COOH ), etc., are some examples of conjugate acid-base pairs.

Explanation:

To get an idea about conjugate acid-base pair, let us consider the ionization of CH₃COOH in an aqueous solution:

CH₃CO₂H(aq) + H₂O(l) ⇌CH₃CO₂ (aq) + H₃O+(aq)

Since CH3COOH is a weak acid, it undergoes partial ionization in the solution, and the above equilibrium is tints established. In the forward reaction, the CH₃COOH molecule donates a proton (H+ion) which is accepted by the H₂O molecule.

Therefore, according to the Bronsted-Lowry concept, CH₂COOH is an acid and H₂O is a base. In the reverse reaction, the H₃O⁺ ion donates a proton which is accepted by a CH₃COO^– ion. Therefore, in the reverse reaction, H₃O⁺ ion acts as an acid and CH₃COO^– as a base.

⇒ \(\underset{\text { acid }}{\mathrm{CH}_3 \mathrm{CO}_2 \mathrm{H}}(a q)+\underset{\text { base }}{\mathrm{H}_2 \mathrm{O}}(l) \underset{\text { base }}{\mathrm{CH}_3 \mathrm{CO}_2^{-}(a q)}+\underset{3}{\mathrm{H}_3 \mathrm{O}^{+}}(a q)\) ………………….(1)

In equation (1), CH₃COO^– ion is the conjugate base of CH₆COOH and CH₃COOH is the conjugate acid of CH₃COO^– ion. Hence, (CH₃COOH, and CH₃COO^–) constitute a conjugate acid-base pair.

Similarly, in equation (1), H₂O is the conjugate base of the H₃O⁺ ion and the H₃O⁺ion is the conjugate acid of H₂O. Therefore, (H₃O⁺, H₂O ) constitutes a conjugate acid-base pair.

An acid donates a proton to produce a conjugate base and a base accepts a proton to produce a conjugate acid. The conjugate base of an acid has one fewer proton than the acid. On the other hand, the conjugate acid of the base has one more proton than the base

Acid -conjugate base:

Base-conjugate base:

Strength of conjugate acid-base pair or Bronsted acid-base pair in aqueous solution:

The stronger an acid, the greater its ability to donate a proton. Similarly, a base with greater proton-accepting ability exhibits stronger basicity.

• The acids HCL, HNO₃, H₂SO₄, etc., undergo complete ionization in aqueous solution to form H₃O⁺ ions and the corresponding conjugate bases. Hence, these are considered strong acids in an aqueous solution. In an aqueous solution, the conjugate base produced from a strong acid has less tendency than H₂O to accept a proton. Therefore in an aqueous solution, the conjugate base of a strong acid is found to be very weak.
• The acids HF, HCN, CH₆COOH, HCOOH, etc., undergo slight ionization in an aqueous solution to produce H₂O+ions and the corresponding conjugate bases. As these acids have little tendency to donate protons in aqueous solution, they are called weak acids.
• In an aqueous solution, the conjugate base produced from a weak acid has more tendency’ than H₂O to accept a proton. Therefore, in aqueous solution, the conjugate base of a weak add is found to be stronger than H₂O.
• In aqueous solution, strong bases like NH₂, O₂^–, H⁺ etc., react completely with water to form their corresponding conjugate adds and OH^– ions. These conjugate acids are later than H₂O. Hence, in aqueous solution, the conjugate acid of a strong base is very weak.

On the other hand, in an aqueous solution, weak bases like NH₃, CH₃NH₂, etc., react partially with water to produce the corresponding conjugate acids and OH^– ions. These conjugate acids are stronger than H₂O. Therefore, in an aqueous solution, the conjugate acid of a weak base is strong.

The conjugate acid of a strong base has little tendency to accept protons. On the other hand, the conjugate acid of a weak base has a high tendency to accept protons.

Acid-base neutralization reactions according to Bronstedlowry concept:

According to the Bronsted-Lowry concept, in an acid-base neutralization reaction, a proton from an acid molecule gets transferred to a molecule of a base.

As a result, the acid converts to its conjugate base, and the base changes to its conjugate acid by accepting a proton.

Example:

Limitations of Bronsted-Lowry concept:

With the help of this theory, the reaction of an acid with a base is explained in terms of the gain or loss of proton(s). However, there are many acid-base reactions in which the exchange of proton(s) does not take place. Such types of acid-base reactions cannot be explained by this theory.

The acidic properties of many non-metallic oxides (for example; CO₂, SO₂ ) and basic properties of many metallic oxides (for example; CaO, BaO) cannot be explained with the help of the Bronsted-Lowry concept. Also, the acidic properties of BF₃, AlCl₂, SnCl₂, etc., cannot be explained with the help of this theory.

Lewis’s Concept Of Acids And Bases

On the electronic theory of valency, scientist Gilbert N. Lewis proposed the following definitions of acids and bases.

Acid: An acid is a substance which can accept a pair of electrons

Examples:

The compounds that have a central atom with incomplete octets can act as Lewis acids, such as BF₃ BCl₃, AlCl₃, etc.

In some compounds due to the presence of vacant orbitals in the central atom, the octet can be expanded. These compounds can also behave as Lewis acids, such as PCl₃, SnCl₃, SiF₄, etc.

SiF₄ (Lewis acid) + 2F^– (Lewis base) → [SiF₆]^2-

Cations like H⁺, Ag⁺, Cu²⁺, Fe³⁺, Al³⁺, etc., behave as Lewis acids.

Molecules containing multiple bonds between two atoms of different electronegativities behave as Lewis acids, such as CO₂, SO₂, etc.

Base: A base Is a Substance that can donate a pair of electrons

Example: Compounds that contain an atom having one or more lone pairs of electrons behave as Lewis bases, such as NH₃, H₂O: CH₃OH, etc.

Anions like NH^–, Cl^–, I^–, OH^–, CN^– etc., are considered as Lewis bases.

• A Lewis add is an acceptor of a pair of electrons and forms a coordinate bond with a Lewis base.
• A Lewis base is a donor of a pair of electrons and forms a coordinate bond with Lewis acid.

Limitations of Lewis’s concept:

• This concept provides no idea regarding the relative strengths of acids and bases.
• This theory contradicts the general concept of acids by considering BF₃ AlCl₃, and simple cations as acids.
• The behaviour ofprotonic adds such as HCl, H₂SO₄ etc., cannot be explained by this concept. These acids do not form coordinate bonds with bases which is the primary requirement Lewis concept.
• Normally, the formation of coordination compounds is slow, therefore acid-base reactions should also be slow, but acid-base reactions are extremely fast, this cannot be explained by the Lewis concept.