Acids, Bases & Salts - Chemistry Class 11 - NEET PDF Download

Table of Contents
1. Arrhenius theory of Acids & Bases
2. Brønsted-Lowry Theory of Acids and Bases
3. Lewis Theory of Acids & Bases
4. Levelling Effect
5. The pH Scale
6. Dissociation of Acids & Bases
7. Buffer Solutions
8. Salts
9. Conjugate Acid-Base Pairs (CABP)
10. Relative Strength of Acids and Bases
11. Ionic Equilibrium
12. Factors Affecting Degree of Ionisation
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Arrhenius theory of Acids & Bases

According to the Arrhenius concept of acids and bases:

• Acid: A substance that increases the concentration of hydrogen ions, H⁺, on dissolving in water. Greater the number of H⁺ ions produced in the solution, the stronger is the acid.
• Base: A substance that increases the concentration of hydroxide ions, OH^-, on dissolving in water. Greater the number of OH^- ions produced in the aqueous solution, the stronger is the base.

Brønsted-Lowry Theory of Acids and Bases

• A Brønsted-Lowry acid is a proton (H⁺) donor; a Brønsted-Lowry base is a proton acceptor.
• A Brønsted-Lowry acid must contain at least one hydrogen atom, typically attached to an electronegative atom such as oxygen or nitrogen.
• When a Brønsted-Lowry acid donates a proton, it forms its conjugate base. When a Brønsted-Lowry base accepts a proton, it forms its conjugate acid.
• The strength of an acid is measured by its ability to donate protons; the strength of a base is measured by its ability to accept protons.
• There is an inverse relation between the strengths of conjugate pairs: the stronger the acid, the weaker its conjugate base; the stronger the base, the weaker its conjugate acid.

Lewis Theory of Acids & Bases

Lewis Acid

• Lewis acids are species that can accept an electron pair; they are electron-pair acceptors.
• Typical Lewis acids are electron-deficient species with empty orbitals (for example, trigonal planar species with an empty p-orbital such as BR₃).
• Species such as H⁺ and many metal cations (Fe³⁺, Mg²⁺, Li⁺, etc.) act as Lewis acids by accepting electron pairs from ligands or bases.
• Some substances (for example, water) can act as both Lewis acids and Lewis bases depending on the reaction partner.
• All Brønsted-Lowry acids are Lewis acids, but the converse need not be true: a Lewis acid need not donate a proton.

Examples of Lewis Acids

1. H⁺ (proton) and onium ions like H₃O⁺.
2. High-oxidation-state transition metal cations such as Fe³⁺.
3. Metal cations that form aqua complexes (Mg²⁺, Li⁺) where water donates electron pairs as ligands.
4. Carbocations such as CH₃⁺ (trigonal planar) which accept electron pairs.
5. Electron-deficient molecular species such as PCl₅, AlCl₃, BF₃ act as Lewis acids.
6. Electron-deficient π-systems (for example, enones) that can accept electron density.

Lewis Base

• Lewis bases are electron-pair donors; they have a filled orbital or a lone pair available for donation.
• Common Lewis bases include ammonia, alkyl amines and other molecules with lone pairs (O, S, N atoms with lone pairs).
• Many Lewis bases are anionic (for example, OH^-, F^-, H^-) and their basic strength often correlates with the acidity (pK_a) of the corresponding conjugate acid.
• Lewis bases are nucleophiles; Lewis acids are electrophiles.

Examples of Lewis Bases

1. Pyridine and its derivatives act as electron-pair donors and behave as Lewis bases.
2. Molecules containing heteroatoms such as O, N or S with lone pairs (e.g., water, alcohols, ketones) act as Lewis bases.
   
3. Simple anions and complex anions that have lone pairs (H^-, F^-, SO₄^2- etc.) can donate electron pairs.
4. Electron-rich π-systems (benzene, ethyne, ethene) can act as Lewis bases in certain reactions.
5. Species with readily available lone pairs such as CH₃^- and OH^- are strong Lewis bases.
6. Weak Lewis acids have relatively strong conjugate Lewis bases.

Limitations of Lewis Concept

1. The Lewis concept is very general but does not by itself explain protonic behaviour (specific proton transfers) of classical acids such as HCl, H₂SO₄, HNO₃.
2. It does not give a quantitative prediction of relative acid-base strengths (pK_a, pK_b values) by itself.

Basic Strength

Basic strength depends primarily on the availability of the lone pair that can accept a proton. Factors affecting availability include inductive effects, resonance (delocalisation), solvation and steric hindrance.

Comparison of Basicity: Ammonia and Alkyl Amines

Compare the basic strength of NH₃, CH₃NH₂, (CH₃)₂NH, (CH₃)₃N

• Factors which affect the basicity of amines: steric effect, inductive effect, and solvation effect.
• The more electron-releasing alkyl groups increase electron density on nitrogen (inductive effect) and generally increase basicity in the gas phase.
• Steric hindrance reduces basic strength by making the lone pair less accessible for protonation.
• Solvation stabilises the conjugate acid. In aqueous solution, solvation effects can reverse gas-phase trends because primary amine conjugate acids form more H-bonds with water than tertiary amine conjugate acids.

• Considering only electronic (inductive) effects in the gas phase, basicity order is (CH₃)₃N > (CH₃)₂NH > CH₃NH₂ > NH₃.
• In aqueous solution, taking solvation into account, the typical order becomes (CH₃)₂NH > CH₃NH₂ > (CH₃)₃N > NH₃.
• Aromatic amines (anilines) are significantly less basic because the lone pair on nitrogen is delocalised (in conjugation with the benzene ring) and less available for protonation.

Levelling Effect

The levelling effect is the observation that a solvent can limit the maximum observable strength of acids and bases. In water, any acid stronger than H₃O⁺ will protonate water to produce H₃O⁺, so in aqueous solution very strong acids (HCl, HNO₃, H₂SO₄) are all effectively levelled to H₃O⁺. Similarly, bases stronger than OH^- are levelled to OH^- in water.

The levelling effect depends on the solvent: in a less ionising solvent (for example, acetic acid) strong acids and bases may display different strengths than in water.

The pH Scale

pH is defined as the negative logarithm of hydrogen ion concentration:

• pH = -log[H⁺]
• [H⁺] = 10^-pH
• pOH = -log[OH^-]
• pH + pOH = 14 (at 25 °C) and Kw = [H⁺][OH^-] = 1.0 × 10^-14 (at 25 °C)

• Very dilute acidic solutions with [H⁺] < 10^-7 M have pH values slightly less than 7 because of the contribution of water ionization; e.g., 10^-8 N HCl solution has pH ≈ 6.958.
• pH is measured accurately by pH meter (electrometric methods) and approximately by indicator papers.
• pH can be zero or negative for very concentrated acids (for example, 1.0 N HCl has pH = 0; more concentrated solutions can have negative pH values).

Dissociation of Acids & Bases

Electrolytes may be strongly or weakly ionised in aqueous solution. Weak electrolytes dissociate only partially; their extent of ionisation is described by the degree of ionisation (α) or by ionisation constants.

For a weak monoprotic acid,

HA + H₂O ⇄ H₃O⁺ + A^-

For a weak base,

B + H₂O ⇄ BH⁺ + OH^-

The equilibrium constants for these reactions are the acid dissociation constant K_a and the base dissociation constant K_b.

Acid and base ionization constants

Buffer Solutions

Definition: A buffer solution is one which resists change in pH when small amounts of acid or base are added.

Acidic buffer: pH < 7. Example: CH₃COOH/CH₃COONa; boric acid/borax.

Basic buffer: pH > 7. Example: NH₄OH/NH₄Cl.

Buffer system in blood: The primary buffer that regulates blood pH is the carbonic acid-bicarbonate system: H₂CO₃/NaHCO₃ (involving dissolved CO₂, H₂CO₃ and HCO₃^-).

Henderson-Hasselbalch Equation

The Henderson-Hasselbalch equation gives an approximate pH of a buffer containing a weak acid and its conjugate base.

For an acidic buffer (weak acid HA and its conjugate base A^-):

pH = pK_a + log([A^-]/[HA])

For a basic buffer (weak base B and its conjugate acid BH⁺):

pOH = pK_b + log([BH⁺]/[B])

Convert to pH using pH = 14 - pOH. Using pK_b = 14 - pK_a (at 25 °C), one can relate buffer pH to pK_a as well.

• pK_a = -log K_a
• pK_b = -log K_b
• K_a and K_b are the dissociation constants of the acid and base respectively.
• [Salt], [Acid] and [Base] represent molar concentrations of the conjugate salt and the weak acid or weak base.

Henderson-Hasselbalch Applied Forms

For an acidic buffer mixture (weak acid HA + conjugate base A^-):

• pK_a = -log K_a.
• [Salt] represents the concentration of the conjugate base (for example CH₃COO^- in CH₃COOH/CH₃COONa).
• [Acid] represents the concentration of the weak acid (HA).

Basic buffer mixture

pOH form:

Converting to pH:

• pH + pOH = 14
• pH = 14 - pOH
• Using pK_b = 14 - pK_a (at 25 °C), the pH of a basic buffer can be related to pK_a as well.

Example: In the buffer NH₄OH/NH₄Cl, NH₄⁺ is the conjugate acid of NH₃. The relevant equilibrium and ionization constant are:

Importance of Buffer Solutions

• Biological processes: Physiological pH (for example, blood pH ≈ 7.4) is maintained by buffer systems such as H₂CO₃/HCO₃^-.
• Industrial processes: Buffers are used in electroplating, manufacture of leather, dyes and photographic materials.
• Analytical chemistry: Buffers are used to remove interfering ions, in complexometric titrations, and to calibrate pH meters.
• Bacteriological research: Culture media are often buffered to keep pH optimal for microbial growth.

Salts

Definition: Salts are ionic compounds formed by the neutralisation reaction between an acid and a base.

Types of Salts

1. Normal salts - Obtained by complete neutralisation of an acid with a base; examples: NaCl, K₂SO₄.
2. Acidic salts - Formed by incomplete neutralisation of polybasic acids; examples: NaHCO₃, NaHSO₄.
3. Basic salts - Formed by incomplete neutralisation of polybasic bases; examples: Mg(OH)Cl, Bi(OH)₂Cl.
4. Double salts - Formed by combination of two different salts and exist typically in the solid state (they dissociate into individual ions in solution); examples: Mohr's salt (ferrous ammonium sulfate, FeSO₄·(NH₄)₂SO₄·6H₂O), alums.
5. Complex salts - Consist of complex ions; they are stable in both the solid state and in solution (for example, K₃[Fe(CN)₆]).
6. Mixed salts - Furnish more than one cation or more than one anion on dissolution; examples: NaKSO₄, CaNa₂(CO₃)₂ (mixed carbonate examples).

Conjugate Acid-Base Pairs (CABP)

In an acid-base reaction, a conjugate acid-base pair consists of two species that differ by a single proton (H⁺). When an acid donates a proton, it forms its conjugate base; when a base accepts a proton, it forms its conjugate acid. This idea is central to the Brønsted-Lowry concept.

General Reaction

• Acid → H⁺ + Conjugate base
• Base + H⁺ → Conjugate acid

Example:

Consider the proton transfer between HCl and NH₃:

• HCl (acid) donates a proton to form Cl^-, its conjugate base.
• NH₃ (base) accepts a proton to form NH₄⁺, its conjugate acid.

Key Point

• A conjugate acid-base pair differs by exactly one proton.
• Example: H₂SO₄ and HSO₄^- are a conjugate pair (differ by one H⁺), whereas H₂SO₄ and SO₄^2- differ by two protons and are not a single conjugate pair.

Relative Strength of Acids and Bases

In acid-base chemistry, any species and its conjugate species have opposite strengths: a strong acid has a weak conjugate base and a strong base has a weak conjugate acid. This relationship helps predict the direction of proton transfer reactions and equilibrium positions.

Example: Strength Order of Some Acids

The input lists a strength order illustrated by an image; the conjugate bases of stronger acids are weaker. See the image for a comparative illustration.

Strength Order of Conjugate Bases

The conjugate bases of acids have the opposite strength order compared to their parent acids; reference figure is provided.

Ionic Equilibrium

An ionic equilibrium in solution is the dynamic balance between unionized molecules and ions formed by ionization. Ionic equilibria determine pH, conductivity and many chemical behaviours of solutions.

Important equilibrium constants (K_eq)

1. Self-ionisation of water

The ionic product of water, Kw = [H⁺][OH^-] = 1.0 × 10^-14 (at 25 °C).

2. Acid dissociation constant (K_a)

For HA ⇄ H⁺ + A^-:

3. Base dissociation constant (K_b)

For B + H₂O ⇄ BH⁺ + OH^-:

Relation between K_a and K_b for a conjugate pair:

• K_a × K_b = K_w (for a conjugate acid-base pair in water).

4. Salt hydrolysis

Some salts react with water to produce acidic or basic solutions. For example, NH₄⁺ from NH₄Cl hydrolyses to produce acidity:

The hydrolysis equilibrium and its constant (K_h) describe this behaviour.

5. Sparingly soluble salts and solubility product (K_sp)

For a salt AB (s) ⇄  A⁺ + B^- the solubility product is:

Factors Affecting Degree of Ionisation

The degree of ionisation (α) of an electrolyte depends on several factors:

• Temperature: Generally, degree of dissociation increases with temperature for many electrolytes; α increases with temperature.
• Dilution: Degree of ionisation increases on dilution. As concentration decreases, dissociation increases (at infinite dilution the limiting behaviour is reached).
• Concentration: Degree of dissociation is inversely related to concentration of the solution; α ∝ 1/[solute] for many weak electrolytes under similar conditions. 
• Nature of solvent: A solvent with a higher dielectric constant weakens attraction between oppositely charged ions and increases ionisation. Water, with a high dielectric constant, is a powerful ionising solvent.
• Presence of common ion: Addition of a strong electrolyte sharing a common ion suppresses ionisation (common-ion effect). For example, ionisation of CH₃COOH is suppressed in presence of a strong source of H⁺.
• Nature of electrolyte: Different electrolytes ionise to different extents even at the same concentration and temperature.

These factors govern equilibria in solutions and are essential for predicting pH, conductance and the outcome of acid-base reactions.