Chemistry Detailed Notes Isc. class 12 CISCE 2026 Syllabus (Page 4)

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1.Solutions

Concentration of Solutions - ISC Class 12 Chemistry

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Study of Concentration of Solutions

1. Concentration of Solutions

It refers to the amount of solute dissolved in a given quantity of solvent or solution.

Common units: Molarity (M), Molality (m), Normality (N), Mole Fraction (χ), Mass %, Volume %

2. Solids in Liquids

Examples: Sugar or salt in water

Solubility: Increases with temperature for most solids.

Concentration: Given by:

\[
\text{Molarity (M)} = \frac{\text{moles of solute}}{\text{volume of solution in L}}
\]

3. Liquids in Liquids

Types: Miscible (ethanol-water), Immiscible (oil-water)

Ideal solution: Obeys Raoult’s Law:

\[
P_{\text{total}} = P_A + P_B = \chi_A P_A^0 + \chi_B P_B^0
\]

Where \( \chi \) = mole fraction and \( P^0 \) = vapor pressure of pure component.

4. Gases in Liquids

Example: CO₂ in soft drinks

Solubility law: Henry’s Law:

\[
C = k_P \cdot P
\]

Where \( C \) = concentration of gas, \( P \) = partial pressure, \( k_P \) = Henry’s law constant.

Solubility decreases with increase in temperature.

5. Solid Solutions

Definition: A solid solution is a homogeneous mixture of two or more solids.

Examples: Alloys (Brass, Bronze, Steel)

Can be substitutional or interstitial depending on the size and type of atoms involved.

Colligative Properties - ISC Class 12 Chemistry

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Colligative Properties

1. Raoult's Law (Relative Lowering of Vapour Pressure)

1st Law: The partial vapour pressure of each component of a solution is directly proportional to its mole fraction.

\[
P_A = \chi_A \cdot P_A^0
\]

Where:

\( P_A \): partial vapour pressure of component A
\( \chi_A \): mole fraction of A
\( P_A^0 \): vapour pressure of pure A

2nd Law (Relative Lowering): The relative lowering of vapour pressure is equal to the mole fraction of the solute.

\[
\frac{P^0 - P}{P^0} = \chi_{\text{solute}} = \frac{n_B}{n_A + n_B}
\]

2. Elevation of Boiling Point

Boiling point of a solution is higher than that of pure solvent.

\[
\Delta T_b = K_b \cdot m
\]

Where:

\( \Delta T_b \): elevation in boiling point
\( K_b \): ebullioscopic constant (boiling point elevation constant)
\( m \): molality of the solution

3. Depression of Freezing Point

Freezing point of a solution is lower than that of the pure solvent.

\[
\Delta T_f = K_f \cdot m
\]

Where:

\( \Delta T_f \): depression in freezing point
\( K_f \): cryoscopic constant (freezing point depression constant)
\( m \): molality of the solution

4. Osmotic Pressure

Osmotic pressure is the pressure required to stop the flow of solvent through a semipermeable membrane into the solution.

\[
\Pi = C R T
\]

Where:

\( \Pi \): osmotic pressure
\( C \): molar concentration (mol/L)
\( R \): universal gas constant
\( T \): temperature in Kelvin

Use of Colligative Properties - ISC Chemistry

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Use of Colligative Properties in Determining Molecular Masses

1. Determination of Molecular Mass Using Colligative Properties

Colligative properties depend only on the number of solute particles, not their identity.

Molecular mass of a solute can be determined by any colligative property using formulas such as:

Using Boiling Point Elevation or Freezing Point Depression:

\[
M = \frac{1000 \cdot K \cdot w_2}{\Delta T \cdot w_1}
\]

Where:

\( M \): Molar mass of solute
\( K \): Cryoscopic (K_f) or Ebullioscopic (K_b) constant
\( w_2 \): Mass of solute (g)
\( w_1 \): Mass of solvent (g)
\( \Delta T \): Change in temperature

2. Abnormal Molecular Mass

Sometimes, calculated molar mass using colligative properties is different from the actual value. This is called abnormal molecular mass.

Occurs due to:

Association: Molecules combine to form larger units (e.g., acetic acid in benzene).
Dissociation: Molecules split into ions (e.g., NaCl in water).

3. Van’t Hoff Factor (i)

Van’t Hoff introduced a correction factor 'i' to account for association or dissociation.

\[
i = \frac{\text{Observed colligative property}}{\text{Calculated colligative property (ideal)}}
\]

Also expressed as:

\[
i = \frac{\text{Normal molecular mass}}{\text{Abnormal molecular mass}}
\]

For dissociation:

\[
i = 1 + \alpha(n - 1)
\]

Where:

\( \alpha \): degree of dissociation
\( n \): number of particles after dissociation

For association:

\[
i = 1 - \alpha \left(1 - \frac{1}{n}\right)
\]

Where \( n \) is the number of molecules associating to form 1 unit.

4. Corrected Colligative Property Formula

To calculate molar mass in cases of dissociation or association:

\[
M = \frac{1000 \cdot K \cdot w_2}{\Delta T \cdot w_1 \cdot i}
\]

Chemistry Concentration Units - ISC Class 12

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Measures of Concentration

1. Molarity (M)

Definition: Molarity is the number of moles of solute per litre of solution.

\[
M = \frac{n}{V}
\]

Where:

\( n \) = moles of solute
\( V \) = volume of solution in litres

Example: 1 mole of NaCl in 1 litre of solution → 1 M solution.

2. Molality (m)

Definition: Molality is the number of moles of solute per kg of solvent.

\[
m = \frac{n}{W}
\]

Where \( W \) is mass of solvent in kg.

Example: 2 moles of KCl in 1 kg of water → 2 m solution.

3. Normality (N)

Definition: Number of gram equivalents of solute per litre of solution.

\[
N = \frac{\text{Gram equivalent of solute}}{\text{Volume of solution in litres}}
\]

Example: 49 g of H₂SO₄ in 1 L (1 equivalent) → 1 N solution.

4. Mole Fraction (χ)

Definition: Ratio of moles of one component to the total moles in solution.

\[
\chi_A = \frac{n_A}{n_A + n_B}
\]

Example: 2 mol ethanol + 3 mol water

\[
\chi_{\text{ethanol}} = \frac{2}{2+3} = 0.4
\]

5. Parts Per Million (ppm)

Definition: Number of parts of solute per 1 million parts of solution.

\[
\text{ppm} = \frac{\text{Mass of solute}}{\text{Mass of solution}} \times 10^6
\]

Example: 2 mg NaCl in 1 kg water:

\[
\text{ppm} = \frac{2}{1000000} \times 10^6 = 2 \, \text{ppm}
\]

🧪 Sample Problems

Q1: Calculate the molarity of 5.8 g NaCl in 500 mL solution.
Na = 23, Cl = 35.5

Solution:

\[
\text{Molar mass of NaCl} = 23 + 35.5 = 58.5 \, \text{g/mol}
\]
\[
\text{Moles} = \frac{5.8}{58.5} = 0.1
\quad ; \quad V = 0.5 \, L
\]
\[
M = \frac{0.1}{0.5} = 0.2 \, \text{M}
\]

Q2: What is the molality of 18 g glucose (C₆H₁₂O₆) in 200 g of water?

\[
\text{Molar mass of glucose} = 180 \, \text{g/mol}
\quad ; \quad n = \frac{18}{180} = 0.1
\]
\[
W = 200\, \text{g} = 0.2\, \text{kg}
\quad ; \quad m = \frac{0.1}{0.2} = 0.5 \, \text{mol/kg}
\]

Henry’s Law – Solubility of Gases in Liquids

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Henry’s Law – Solubility of Gases in Liquids

1. Statement of Henry's Law

Henry’s Law: The mass (or solubility) of a gas dissolved in a liquid is directly proportional to the pressure of the gas above the liquid at constant temperature.

\[
C = k_H \cdot P
\]

Where:

\( C \) = concentration or solubility of the gas (mol/L)
\( P \) = partial pressure of the gas
\( k_H \) = Henry’s law constant (depends on gas and temperature)

2. Applications of Henry's Law

Bottling of soft drinks and soda (CO₂ is dissolved under high pressure)
Scuba diving (oxygen and nitrogen dissolve in blood under pressure)
Decompression sickness (gas bubbles form if pressure drops too fast)

3. Simple Numerical Problems

Example 1:

Q: At 298 K, Henry’s law constant for oxygen is \( k_H = 1.3 \times 10^5 \, \text{atm} \). Calculate the solubility of oxygen in water if the partial pressure is 0.21 atm.

Solution:

\[
C = \frac{P}{k_H} = \frac{0.21}{1.3 \times 10^5} \approx 1.615 \times 10^{-6} \, \text{mol/L}
\]

Example 2:

Q: The solubility of CO₂ in water at 1 atm is 0.030 mol/L. What will be the solubility at 2 atm?

Solution:

\[
C_1 = 0.030 \, \text{mol/L at } P_1 = 1 \, \text{atm}
\]
\[
C_2 = \frac{P_2}{P_1} \cdot C_1 = \frac{2}{1} \cdot 0.030 = 0.060 \, \text{mol/L}
\]

4. Important Notes

Higher pressure → greater solubility of gas in liquid
Higher temperature → lower solubility (gases escape easily)
Units of \( k_H \) vary: commonly atm, bar, or mol/(L·atm)

Raoult’s Law, Ideal & Non-Ideal Solutions, Azeotropes

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Raoult’s Law and Solutions

1. Raoult’s Law for Non-Volatile Solutes

Raoult’s Law states that the partial vapor pressure of the solvent is directly proportional to its mole fraction in the solution.

\[
P_1 = X_1 P_1^0
\]

\( P_1 \): Vapor pressure of solvent in solution
\( X_1 \): Mole fraction of solvent
\( P_1^0 \): Vapor pressure of pure solvent

The relative lowering of vapor pressure is given by:

\[
\frac{P_1^0 - P_1}{P_1^0} = X_2
\]

\(X_2\): Mole fraction of solute

2. Raoult’s Law for Volatile Solutes

When both solute and solvent are volatile:

\[
P_{\text{total}} = P_1 + P_2 = X_1 P_1^0 + X_2 P_2^0
\]

\(P_{\text{total}}\): Total vapor pressure
\(X_1, X_2\): Mole fractions of components
\(P_1^0, P_2^0\): Pure component vapor pressures

3. Ideal Solution

Obeys Raoult’s Law at all compositions
No enthalpy or volume change on mixing: \( \Delta H_{\text{mix}} = 0 \), \( \Delta V_{\text{mix}} = 0 \)
Examples: benzene & toluene, n-hexane & n-heptane

4. Non-Ideal Solution

Deviation from Raoult’s Law
\( \Delta H_{\text{mix}} \neq 0 \), \( \Delta V_{\text{mix}} \neq 0 \)
Positive deviation: total vapor pressure > expected (e.g., ethanol + acetone)
Negative deviation: total vapor pressure < expected (e.g., acetone + chloroform)

5. Azeotropic Mixtures

Definition: A mixture of two liquids that boils at a constant temperature and behaves like a pure substance.

Types:

Minimum boiling azeotrope: shows positive deviation (e.g., ethanol + water at 95.6%)
Maximum boiling azeotrope: shows negative deviation (e.g., HCl + water at 20.2%)

6. Graphical Representation

Graph of vapor pressure vs mole fraction:

Ideal solution: Straight line between pure components
Non-ideal (positive/negative deviation): Curve lies above/below ideal line

7. Fractional Distillation

Used to separate liquid mixtures with different boiling points
Based on difference in vapor pressures
Limitation: Cannot separate azeotropes by simple distillation
Example: Ethanol-water mixture (forms azeotrope at 95.6% ethanol)

Colligative Properties & Molecular Mass Determination

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Colligative Properties

1. Definition

Colligative properties are properties of dilute solutions that depend only on the number of solute particles (moles), not on their nature or identity.

These include:

Relative lowering of vapor pressure
Elevation of boiling point
Depression of freezing point
Osmotic pressure

2. Examples

Adding salt to water raises the boiling point – elevation of boiling point
Adding ethylene glycol to car radiators – depression of freezing point
Preserving food with sugar/salt – osmotic pressure prevents bacterial growth

3. Determination of Molecular Mass using Colligative Properties

Colligative properties are used to find the molar mass \( M_2 \) of an unknown solute based on changes observed in a known solvent.

Formulas:

(i) Using Depression of Freezing Point:

\[
\Delta T_f = K_f \cdot m = K_f \cdot \frac{W_2 \times 1000}{M_2 \times W_1}
\]

(ii) Using Elevation of Boiling Point:

\[
\Delta T_b = K_b \cdot m = K_b \cdot \frac{W_2 \times 1000}{M_2 \times W_1}
\]

(iii) Using Osmotic Pressure:

\[
\Pi = CRT = \frac{n_2}{V}RT = \frac{W_2 RT}{M_2 V}
\]

\( \Delta T_f, \Delta T_b \): Freezing/Boiling point changes
\( K_f, K_b \): Cryoscopic/Ebullioscopic constant
\( W_2 \): Mass of solute, \( W_1 \): Mass of solvent
\( M_2 \): Molar mass of solute
\( \Pi \): Osmotic pressure, \( C \): Concentration, \( R \): Gas constant, \( T \): Temperature

4. Key Point

The more the depression/elevation/pressure change, the greater the number of solute particles—this allows for indirect determination of unknown molecular masses.

Relative Lowering of Vapour Pressure

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Relative Lowering of Vapour Pressure

1. Definition

The relative lowering of vapour pressure refers to the ratio of the decrease in vapour pressure of a solvent when a non-volatile solute is added, to the vapour pressure of the pure solvent.

2. Raoult’s Law – Mathematical Expression

Raoult’s Law: The partial vapour pressure of each component of a solution is directly proportional to its mole fraction.

For a solution containing a non-volatile solute:

\[
\frac{P^0 - P_s}{P^0} = \frac{n_2}{n_1 + n_2} \approx \frac{n_2}{n_1}
\]

\( P^0 \): Vapour pressure of pure solvent
\( P_s \): Vapour pressure of solution
\( n_1 \): Moles of solvent
\( n_2 \): Moles of solute

When \( n_2 \ll n_1 \), \( n_1 + n_2 \approx n_1 \), so:

\[
\frac{\Delta P}{P^0} = \frac{n_2}{n_1}
\]

3. Determination of Molecular Mass Using Vapour Pressure

We use Raoult’s Law to determine the relative molecular mass of a non-volatile solute:

\[
\frac{\Delta P}{P^0} = \frac{W_2 \cdot M_1}{W_1 \cdot M_2}
\]

\( \Delta P = P^0 - P_s \): Lowering of vapour pressure
\( W_2 \): Mass of solute
\( W_1 \): Mass of solvent
\( M_1 \): Molar mass of solvent
\( M_2 \): Molar mass of solute (to be calculated)

4. Key Applications

Used in determining molecular masses of unknown solutes.
Effective only when the solute is non-volatile and does not dissociate/associate in the solvent.

Depression in Freezing Point

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Depression in Freezing Point

1. Definition

The freezing point of a liquid is lowered when a non-volatile solute is added to it. The phenomenon is known as depression in freezing point.

The difference between the freezing point of the pure solvent and the freezing point of the solution is called depression in freezing point:
\[
\Delta T_f = T_f^0 - T_f
\]
where:

\( T_f^0 \): Freezing point of pure solvent
\( T_f \): Freezing point of solution

2. Cryoscopic Constant (\( K_f \))

The depression in freezing point is directly proportional to the molality of the solution.

\[
\Delta T_f = K_f \cdot m
\]
where:

\( \Delta T_f \): Depression in freezing point
\( K_f \): Cryoscopic constant (molal depression constant)
\( m \): Molality of the solution

3. Mathematical Derivation

Consider:
\[
m = \frac{n_2}{W_1 \ (\text{in kg})} = \frac{W_2}{M_2 \cdot W_1}
\]
Substituting this in the formula:
\[
\Delta T_f = K_f \cdot \frac{W_2}{M_2 \cdot W_1}
\]
Rearranging to calculate molar mass:
\[
M_2 = \frac{K_f \cdot W_2}{\Delta T_f \cdot W_1}
\]

\( W_2 \): Mass of solute (g)
\( M_2 \): Molar mass of solute (g/mol)
\( W_1 \): Mass of solvent (kg)
\( \Delta T_f \): Freezing point depression

4. Units and Notes

The unit of \( K_f \) is:
\[
\text{K} \cdot \text{kg mol}^{-1}
\]
The value of \( K_f \) depends on the nature of the solvent.

Elevation in Boiling Point

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Elevation in Boiling Point

1. Definition

The boiling point of a solution is higher than that of the pure solvent. The increase in boiling point caused by the presence of a non-volatile solute is called the elevation in boiling point.

The elevation in boiling point is defined as:
\[
\Delta T_b = T_b - T_b^0
\]
where:

\( T_b^0 \): Boiling point of pure solvent
\( T_b \): Boiling point of solution

2. Ebullioscopic Constant (\( K_b \))

The elevation in boiling point is directly proportional to the molality of the solution:
\[
\Delta T_b = K_b \cdot m
\]
where:

\( \Delta T_b \): Elevation in boiling point
\( K_b \): Ebullioscopic constant
\( m \): Molality of the solution

3. Mathematical Derivation

Molality is defined as:
\[
m = \frac{n_2}{W_1 \ (\text{in kg})} = \frac{W_2}{M_2 \cdot W_1}
\]
Substituting in the formula:
\[
\Delta T_b = K_b \cdot \frac{W_2}{M_2 \cdot W_1}
\]
Rearranged to calculate molar mass:
\[
M_2 = \frac{K_b \cdot W_2}{\Delta T_b \cdot W_1}
\]

\( W_2 \): Mass of solute (g)
\( M_2 \): Molar mass of solute (g/mol)
\( W_1 \): Mass of solvent (kg)
\( \Delta T_b \): Boiling point elevation

4. Units and Notes

The unit of \( K_b \) is:
\[
\text{K} \cdot \text{kg mol}^{-1}
\]
The value of \( K_b \) is constant for a given solvent.

Osmotic Pressure Notes

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Osmotic Pressure and Related Concepts

1. Osmotic Pressure: Definition & Explanation

Osmotic pressure is the pressure required to stop the flow of solvent molecules through a semipermeable membrane from a dilute solution into a concentrated one. It is a colligative property depending on solute concentration and temperature.

2. Semipermeable Membranes

Natural Semipermeable Membranes: Found in biological systems (e.g., cell membranes).
Chemical Semipermeable Membranes: Artificial membranes made from materials like cellulose acetate used in reverse osmosis.

3. Reverse Osmosis

A process where pressure is applied to force solvent through a semipermeable membrane from a concentrated solution to a dilute one—used in water purification.

4. Tonicity of Solutions

Isotonic: Solutions with equal osmotic pressure. No net water movement.
Hypotonic: Lower solute concentration than the cell. Water enters the cell.
Hypertonic: Higher solute concentration than the cell. Water leaves the cell.

5. Diffusion vs Osmosis

Diffusion: Movement of particles from high to low concentration without a membrane.
Osmosis: Movement of water through a semipermeable membrane toward higher solute concentration.

6. Applications of Osmotic Pressure

Used to determine the relative molecular mass of solutes, especially macromolecules. The formula is:

π = CRT

π = Osmotic pressure
C = Molar concentration
R = Gas constant (0.0821 L·atm·K⁻¹·mol⁻¹)
T = Absolute temperature in Kelvin

7. Van’t Hoff’s Laws

Van’t Hoff–Boyle’s Law: At constant temperature, osmotic pressure ∝ concentration.
Van’t Hoff–Charles’ Law: At constant concentration, osmotic pressure ∝ absolute temperature.
Van’t Hoff–Avogadro’s Law: Equal volumes of solutions with equal osmotic pressure contain equal number of solute particles.

Abnormal Molecular Mass

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Notes on Abnormal Molecular Mass

(e) Abnormal Molecular Mass

The molecular mass calculated using colligative properties sometimes deviates from the expected (normal) value. This deviation is termed abnormal molecular mass. It usually occurs due to either association or dissociation of solute particles in the solution.

1. Dissociation

When a solute dissociates into two or more particles in solution, the number of particles increases, resulting in a lower calculated molecular mass than expected.

Van’t Hoff factor (i) = measured colligative property / calculated colligative property

i > 1 for dissociation
Example: NaCl in water → Na⁺ + Cl⁻
Expected molar mass of NaCl = 58.5 g/mol, but due to dissociation, it appears lower.

2. Association

When two or more solute particles combine (associate) to form larger molecules, the number of solute particles decreases, leading to a higher molecular mass than expected.

i < 1 for association
Example: Acetic acid in benzene associates as dimer: 2 CH₃COOH ⇌ (CH₃COOH)₂
Apparent molar mass becomes nearly double the actual value.

3. Van’t Hoff Factor (i)

It quantifies the effect of association or dissociation:

i = Normal molar mass / Observed molar mass

Or based on degree of dissociation (α):

i = 1 + (n - 1)α, where n = number of particles formed on dissociation

Example: For Na₂SO₄ ⇌ 2Na⁺ + SO₄²⁻, n = 3
i = 1 + 2α

van’t Hoff Factor and Colligative Properties

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Notes: Van’t Hoff Factor (i) & Colligative Properties

1. Van’t Hoff Factor (i)

The van’t Hoff factor is the ratio of the actual number of particles in solution after dissociation or association to the number of formula units initially dissolved.

i = \(\dfrac{\text{Observed colligative property}}{\text{Calculated colligative property assuming no dissociation/association}}\)

i > 1: for dissociating solutes (electrolytes like NaCl)
i < 1: for associating solutes (e.g. acetic acid in benzene)
i = 1: for nonelectrolytes (e.g. glucose)

2. Modified Colligative Property Formulas

The colligative property formulas are modified by multiplying with i:

ΔT_b = i × K_b × m
ΔT_f = i × K_f × m
π = i × C × R × T
ΔP = i × X_solute × P⁰_solvent

3. Degree of Dissociation (α)

If a solute dissociates into n particles, then:

i = 1 + (n − 1)α

Example: For NaCl → Na⁺ + Cl⁻ (n = 2),
i = 1 + (2 − 1)α = 1 + α

4. Degree of Association (α)

If n molecules associate to form 1 unit:

i = 1 − α(n − 1)/n

Example: 2 acetic acid molecules form a dimer (n = 2),
i = 1 − α(2 − 1)/2 = 1 − α/2

5. Sample Problems

Q: A 0.1 M NaCl solution shows an osmotic pressure of 4.92 atm at 298 K. Calculate the van’t Hoff factor. (R = 0.0821 L·atm/mol·K)

Solution:
π = iCRT → 4.92 = i × 0.1 × 0.0821 × 298
i = 4.92 / (0.1 × 0.0821 × 298) ≈ 2.0
➤ Almost complete dissociation.

Q: The observed molar mass of acetic acid in benzene is 120 g/mol. The normal molar mass is 60 g/mol. Calculate the van’t Hoff factor and degree of association.

i = 60 / 120 = 0.5
For dimerization (n = 2), i = 1 − α/2 → 0.5 = 1 − α/2
α = 1.0 → 100% association into dimers.

Colligative Properties: Numerical Problems

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Numerical Problems – Colligative Properties

1. Osmotic Pressure

Q1: Calculate the osmotic pressure of a 0.1 M solution of glucose at 298 K. (R = 0.0821 L·atm·K⁻¹·mol⁻¹)

π = C × R × T

π = 0.1 × 0.0821 × 298 ≈ 2.45 atm

Q2: A solution of NaCl (i = 2) has a concentration of 0.2 M at 300 K. Find its osmotic pressure.

π = iCRT = 2 × 0.2 × 0.0821 × 300 ≈ 9.85 atm

2. Abnormal Molecular Mass

Q3: The observed molar mass of a solute is 30 g/mol, but its normal molar mass is 60 g/mol. Is this due to dissociation or association?

i = Normal / Observed = 60 / 30 = 2

Since i > 1, the solute undergoes dissociation.

Q4: The observed molar mass of acetic acid in benzene is 120 g/mol (actual = 60 g/mol). Determine the van’t Hoff factor.

i = 60 / 120 = 0.5

Since i < 1, the solute undergoes association.

3. Van’t Hoff Factor & Degree of Dissociation/Association

Q5: If Na₂SO₄ dissociates as Na₂SO₄ → 2Na⁺ + SO₄²⁻ and i = 2.5, find the degree of dissociation (α).

i = 1 + (n − 1)α → 2.5 = 1 + 2α → α = 0.75

Degree of dissociation is 75%.

Q6: CH₃COOH dimerizes in benzene (n = 2), i = 0.6. Calculate degree of association (α).

i = 1 − α(n − 1)/n → 0.6 = 1 − α(1)/2 → α = 0.8

Degree of association is 80%.

4. Colligative Property Adjustments

Q7: Calculate the boiling point elevation for a 0.1 m NaCl solution (i = 2), K_b = 0.52 °C·kg/mol.

ΔT_b = i × K_b × m = 2 × 0.52 × 0.1 = 0.104 °C

Q8: A 0.2 mol/kg solution of glucose (non-electrolyte) freezes at what temperature? (K_f = 1.86 °C·kg/mol)

ΔT_f = i × K_f × m = 1 × 1.86 × 0.2 = 0.372 °C

Freezing point = 0 − 0.372 = −0.372 °C

Solutions Chapter – Numerical Problems

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Numerical Problems – Chapter: Solutions

1. Mole Fraction & Molality

Q1: Calculate the mole fraction of NaOH in a solution containing 40 g of NaOH in 180 g of water.

Moles of NaOH = 40 / 40 = 1 mol
Moles of H₂O = 180 / 18 = 10 mol

χ_NaOH = 1 / (1 + 10) = 1/11 ≈ 0.0909

Q2: Calculate the molality of a solution containing 5.8 g of KCl in 200 g of water.

Moles of KCl = 5.8 / 74.5 ≈ 0.0779 mol
Mass of water = 200 g = 0.2 kg

Molality = 0.0779 / 0.2 = 0.3895 mol/kg

2. Abnormal Molecular Mass & Van’t Hoff Factor

Q3: A 0.1 molal solution of KCl freezes at −0.372 °C. Calculate the van’t Hoff factor. (K_f = 1.86 °C·kg/mol)

ΔT_f = i × K_f × m → 0.372 = i × 1.86 × 0.1

i = 0.372 / (1.86 × 0.1) = 2.0 (full dissociation)

Q4: The molar mass of acetic acid in benzene is found to be 120 g/mol. Calculate the van’t Hoff factor. (Normal molar mass = 60 g/mol)

i = Normal / Observed = 60 / 120 = 0.5 (indicates dimerization)

3. Colligative Properties

Q5: Calculate the osmotic pressure of a 0.2 M urea solution at 298 K. (R = 0.0821 L·atm/mol·K)

π = iCRT = 1 × 0.2 × 0.0821 × 298 = 4.89 atm

Q6: Find the elevation in boiling point of a 1 molal NaCl solution. (i = 2, K_b = 0.52 °C·kg/mol)

ΔT_b = i × K_b × m = 2 × 0.52 × 1 = 1.04 °C

Q7: What is the freezing point of a solution of 0.1 mol glucose in 1 kg water? (K_f = 1.86 °C·kg/mol)

ΔT_f = i × K_f × m = 1 × 1.86 × 0.1 = 0.186 °C
T_f = 0 − 0.186 = −0.186 °C

2. Electrochemistry

Electrochemistry Notes

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Electrochemistry – HTML Notes with Math

1. Electrolytic vs Electrochemical Cells

Electrolytic Cell: Uses electrical energy to drive a non-spontaneous redox reaction (e.g., electrolysis of water).
Electrochemical (Galvanic) Cell: Converts chemical energy into electrical energy via spontaneous redox reactions (e.g., Daniell Cell).

2. Redox Reactions in Electrochemical Cells

Oxidation: Loss of electrons (at the anode)

Reduction: Gain of electrons (at the cathode)

Electron flow: From anode → cathode

Example: In the Daniell cell:
Zn(s) → Zn²⁺ + 2e⁻ (oxidation at anode)
Cu²⁺ + 2e⁻ → Cu(s) (reduction at cathode)

3. Electromotive Force (EMF) of a Cell

EMF is the potential difference between two electrodes when no current is flowing.

EMF = E_cathode − E_anode

Measured in volts (V)

4. Standard Electrode Potential (E°)

The potential of a half-cell under standard conditions (1 M, 1 atm, 25°C).

Standard hydrogen electrode (SHE) is taken as reference with E° = 0 V
More positive E° = greater tendency to get reduced

5. Nernst Equation

Relates electrode potential to concentration/pressure of species involved.

E = E° − (0.0591 / n) log Q

Where:

E: Electrode potential
E°: Standard electrode potential
n: Number of electrons transferred
Q: Reaction quotient

Example: For the reaction Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s),
E = E° − (0.0591 / 2) log([Zn²⁺]/[Cu²⁺])

6. Gibbs Free Energy and EMF

Relates cell EMF to the spontaneity of the reaction.

ΔG = −nFE

ΔG: Gibbs free energy (Joules)
n: Number of electrons transferred
F: Faraday constant = 96,500 C/mol
E: EMF of the cell (Volts)

If ΔG < 0 → reaction is spontaneous (E > 0)

Conductance & Electrolysis Notes

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Conductance, Electrolysis & Electrochemical Devices

1. Conductance in Electrolytic Solutions

Conductance (G) = 1 / Resistance (R), unit: S (Siemens)
Measured using a conductivity cell and Wheatstone bridge

2. Types of Conductivity

Specific Conductivity (κ): Conductance of 1 cm³ of solution
κ = G × l/A
Unit: S·cm⁻¹
Molar Conductivity (Λ_m):
Λ_m = κ / C
where C is concentration in mol·L⁻¹
Equivalent Conductivity (Λ_eq):
Λ_eq = κ / N
N = normality

3. Variation with Concentration

Trend:

κ decreases with dilution (fewer ions per volume)
Λ_m increases with dilution (more ion mobility)
Strong electrolytes: smooth increase
Weak electrolytes: sharp increase (ionization increases)

4. Kohlrausch’s Law

Statement: At infinite dilution, molar conductivity is the sum of individual ion conductivities:

Λ^∞_m

Electrochemical Cells & Redox Reactions

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Electrochemical Cells & Redox Reactions

1. Introduction to Electrochemical Cells

An electrochemical cell is a device that converts chemical energy into electrical energy using spontaneous redox reactions.

It consists of two half-cells, each containing an electrode and an electrolyte. The two half-cells are connected via a salt bridge that allows ion exchange and maintains electrical neutrality.

Types:

Galvanic (Voltaic) Cell: Produces electricity from a spontaneous redox reaction
Electrolytic Cell: Uses electricity to drive a non-spontaneous reaction

2. Principle of Redox Reactions

Redox stands for reduction-oxidation reactions, where electrons are transferred:

Oxidation: Loss of electrons (occurs at the anode)
Reduction: Gain of electrons (occurs at the cathode)

The flow of electrons occurs from the anode to the cathode through the external circuit.

Oxidation Half Reaction: Zn(s) → Zn²⁺ + 2e⁻
Reduction Half Reaction: Cu²⁺ + 2e⁻ → Cu(s)

Example – Daniell Cell:

Anode: Zn(s) | Zn²⁺ (Oxidation)
Cathode: Cu²⁺ | Cu(s) (Reduction)
Overall reaction: Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s)
Electron flow: Zn → Cu (anode to cathode)

3. Key Takeaways

Oxidation = Anode (Always)
Reduction = Cathode (Always)
Electrons flow from Anode → Cathode
Ions flow through Salt Bridge
Redox reactions drive the cell

Galvanic Cells - Electrochemistry

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Galvanic Cells (Voltaic Cells)

1. Introduction

A Galvanic cell is an electrochemical cell that converts chemical energy into electrical energy through a spontaneous redox reaction.

It consists of two half-cells connected by a salt bridge and an external wire for electron flow.

2. Representation of a Galvanic Cell

A Galvanic cell is written in the following format:

Anode | Anode electrolyte || Cathode electrolyte | Cathode

Double line (||) represents the salt bridge.

Example – Daniell Cell:

Zn(s) | Zn²⁺(aq) || Cu²⁺(aq) | Cu(s)

3. Principle – Redox Reaction

Galvanic cells are based on redox reactions:

Oxidation: Occurs at the anode (loss of electrons)
Reduction: Occurs at the cathode (gain of electrons)

Zn(s) → Zn²⁺(aq) + 2e⁻ (Oxidation)
Cu²⁺(aq) + 2e⁻ → Cu(s) (Reduction)

4. Mechanism of Production of Electric Current

At the Anode: Metal loses electrons (oxidation). These electrons travel through the external wire.
Electron Flow: Electrons flow from anode → cathode through a metallic wire.
At the Cathode: Positive ions in solution gain electrons (reduction) and get deposited.
Salt Bridge: Maintains charge balance by allowing ion migration (anions to anode, cations to cathode).

      Current flows in the external circuit from cathode → anode (opposite to electron flow).

      Electron flow: anode → cathode

      Ion flow: via salt bridge

5. Summary

Anode: Oxidation (electron release)
Cathode: Reduction (electron acceptance)
Electron flow: Anode → Cathode
Salt bridge maintains neutrality

Electrode Potentials & SHE

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Electrode Potentials & Standard Hydrogen Electrode (SHE)

1. Measurement of Electrode Potential

The electrode potential is the tendency of an electrode to lose or gain electrons when it is in contact with its own ions.

Single Electrode Potential (SEP): Measured by connecting the electrode to the Standard Hydrogen Electrode (SHE).

2. Standard Hydrogen Electrode (SHE)

Definition: SHE is a reference electrode with a potential of 0.00 V under standard conditions.

Construction:

Platinum electrode coated with platinum black
1M H⁺ ion solution (usually HCl)
H₂ gas bubbled at 1 atm over electrode

SHE half-cell reaction: H⁺(aq) + e⁻ ⇌ ½ H₂(g)

Applications: Used as a reference for measuring standard electrode potentials of other half-cells.
Limitations:

Difficult to maintain 1 atm H₂ pressure
Platinum surface may get poisoned

3. Standard Electrode Potential (E°)

Definition: Electrode potential measured under standard conditions: 1 M solution, 1 atm gas, 25°C (298 K).

Example:
Zn²⁺ + 2e⁻ ⇌ Zn(s), E° = −0.76 V
Cu²⁺ + 2e⁻ ⇌ Cu(s), E° = +0.34 V

Measured using: SHE as reference.

Cell setup:

Zn | Zn²⁺ (1M) || H⁺ (1M), H₂(1 atm) | Pt (for Zn potential)
Cu | Cu²⁺ (1M) || H⁺ (1M), H₂(1 atm) | Pt (for Cu potential)

4. Cell Notation – Representation

Standard cell notation follows:

Anode | Anode Electrolyte || Cathode Electrolyte | Cathode

Example (Zn-Cu cell):
Zn(s) | Zn²⁺(1M) || Cu²⁺(1M) | Cu(s)

5. Factors Affecting Electrode Potential

Concentration: Increasing ion concentration shifts equilibrium, changes E. (Use Nernst Equation)
Temperature: Affects kinetics and equilibrium constants; electrode potential may increase or decrease.
Nature of Electrode: Different metals have different tendencies to lose/gain electrons due to their atomic structure.

Nernst Equation: E = E° − (0.0591 / n) log([Red]/[Ox])

Electrochemical Series & Reaction Feasibility

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Electrochemical Series & Prediction of Feasibility

1. What is the Electrochemical Series?

The Electrochemical Series is a list of standard reduction potentials (E°) of various elements and ions arranged in decreasing order.

Elements with higher E° values are stronger oxidizing agents (reduced easily).
Elements with lower E° values are stronger reducing agents (oxidized easily).

2. Standard Reduction Potentials (E°)

Measured under standard conditions: 1 M concentration, 1 atm gas pressure, and 25°C (298 K).

Cu²⁺ + 2e⁻ → Cu E° = +0.34 V
Zn²⁺ + 2e⁻ → Zn E° = −0.76 V

Higher E° → greater tendency to gain electrons (reduction)

3. Predicting Feasibility of Redox Reactions

A redox reaction is spontaneous if the overall cell potential (E°_cell) is positive:

E°_cell = E°_cathode − E°_anode

Example:
Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s)
E°_cell = 0.34 V − (−0.76 V) = +1.10 V → Feasible ✅

Cu(s) + Zn²⁺(aq) → Cu²⁺(aq) + Zn(s)
E°_cell = −0.76 − 0.34 = −1.10 V → Not feasible ❌

4. Summary – How to Use the Electrochemical Series

Step 1: Identify oxidation (anode) and reduction (cathode) half-reactions
Step 2: Get their standard E° values from the electrochemical series
Step 3: Calculate E°_cell = E°_cathode − E°_anode
Step 4: If E°_cell > 0 → reaction is spontaneous

Nernst Equation and Free Energy

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Nernst Equation & Free Energy of Reactions

1. Nernst Equation

The Nernst Equation relates the electrode potential under non-standard conditions to standard electrode potential.

E = E° − (0.0591 / n) × log([Products]/[Reactants])

Where:

E: Electrode potential
E°: Standard electrode potential
n: Number of electrons transferred

For the half-cell: Zn²⁺ + 2e⁻ ⇌ Zn
E = E° − (0.0591/2) × log(1/[Zn²⁺])

2. Free Energy and Cell EMF

The free energy change (ΔG) of a redox reaction is related to the EMF of the cell:

ΔG = −nFE

Where:

ΔG: Gibbs free energy (J)
n: Number of electrons
F: Faraday constant (96500 C/mol)
E: EMF of the cell (V)

      If ΔG < 0 → Reaction is spontaneous

      If ΔG > 0 → Non-spontaneous

3. Relationship Between Free Energy & Equilibrium Constant

At equilibrium, the free energy is related to the equilibrium constant (K) as:

ΔG° = −RT ln K

Also:
E° = (0.0591 / n) × log K

4. Prediction of Spontaneity

To determine whether a redox reaction is spontaneous:

Calculate E°_cell = E°_cathode − E°_anode
If E°_cell > 0, then ΔG < 0 and the reaction is spontaneous.

Cu²⁺ + Zn → Cu + Zn²⁺
E°_cell = 0.34 − (−0.76) = +1.10 V → Spontaneous ✅

5. Example Numerical Problems

Q: Calculate ΔG for a cell reaction where E_cell = 1.1 V and n = 2.
Sol: ΔG = −nFE = −2 × 96500 × 1.1 = −212300 J = −212.3 kJ

Q: If E°_cell = 0.295 V and n = 2, find the equilibrium constant K.
Sol:

log K = (n × E°) / 0.0591 = (2 × 0.295) / 0.0591 ≈ 9.98
K ≈ 10⁹.⁹⁸ ≈ 9.55 × 10⁹

Conductance and Resistance - Notes

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AJAX - VI: Conductance and Resistance

1. Comparison of Metallic and Electrolytic Conductance

Metallic Conductance: Due to flow of free electrons.
Electrolytic Conductance: Due to movement of ions in solution or molten state.
Temperature Effect: Metallic conductance decreases with ↑ temperature; Electrolytic conductance increases with ↑ temperature.
Nature: Metals remain unchanged; Electrolytes undergo chemical changes (electrolysis).

2. Relationship Between Conductance and Resistance

Conductance (G) is the reciprocal of resistance (R):

G = 1 / R

Unit of resistance: ohm (Ω)
Unit of conductance: siemens (S) or mho

3. Specific Resistance and Specific Conductance

Specific Resistance (Resistivity, ρ): Resistance of a conductor of unit length and unit cross-sectional area.

ρ = R × (A / l)

Specific Conductance (Conductivity, κ): Conductance of a solution between two electrodes 1 cm apart and 1 sq. cm area.

κ = G × (l / A)

Unit:
- Resistivity (ρ): Ω·cm
- Conductivity (κ): S·cm⁻¹

Conductance and Cell Constant - AJAX VI

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AJAX - VI: Cell Constant & Conductance Concepts

1. Cell Constant

The cell constant (denoted by K) relates the physical dimensions of a conductivity cell:

Cell Constant (K) = l / A

l = distance between electrodes (cm)
A = cross-sectional area of electrodes (cm²)

To calculate: Use a standard solution (like KCl) of known conductivity (κ):

K = κ / G

κ = specific conductance (S·cm⁻¹)
G = conductance (S)

2. Equivalent Conductance (Λ_eq)

It is the conductance of all the ions produced by one gram-equivalent of an electrolyte in a given volume.

Λ_eq = κ × (1000 / N)

κ = specific conductance
N = normality of the solution

Units: S·cm²·equiv⁻¹

3. Molar Conductance (Λ_m)

Molar conductance is the conductance of all ions produced by one mole of electrolyte in solution.

Λ_m = κ × (1000 / M)

κ = specific conductance
M = molarity (mol/L)

Units: S·cm²·mol⁻¹

4. General Relationships

Specific Conductance (κ) gives info per cm³ of solution.
Molar Conductance (Λ_m) increases as solution is diluted.
Equivalent Conductance (Λ_eq) is useful for acids/bases in normality terms.

Λ_m = κ × (1000 / M)
Λ_eq = κ × (1000 / N)

Graphical Representation:

For strong electrolytes: Λ_m vs √C is linear (slightly decreasing).
For weak electrolytes: Λ_m increases sharply on dilution (nonlinear).

Conductance - Units and Numericals (AJAX VI)

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AJAX VI – Units and Numericals in Conductance

1. Important Units

Quantity	Symbol	Unit
Resistance	R	Ohm (Ω)
Conductance	G	Siemens (S)
Specific Conductance	κ (kappa)	S·cm⁻¹
Cell Constant	K	cm⁻¹
Molar Conductance	Λ_m	S·cm²·mol⁻¹
Equivalent Conductance	Λ_eq	S·cm²·equiv⁻¹

2. Key Formulas

G = 1 / R

κ = G × Cell Constant

Λ_m = κ × (1000 / M)

Λ_eq = κ × (1000 / N)

3. Numerical Examples

Example 1: Calculating Specific Conductance (κ)

Given:

Conductance (G) = 0.005 S
Cell constant (K) = 1.2 cm⁻¹

κ = G × K = 0.005 × 1.2 = 0.006 S·cm⁻¹

Example 2: Finding Molar Conductance (Λ_m)

Given:

κ = 0.006 S·cm⁻¹
Molarity (M) = 0.01 mol/L

Λ_m = 0.006 × (1000 / 0.01) = 600 S·cm²·mol⁻¹

Example 3: Calculate Cell Constant

Given:

κ = 0.014 S·cm⁻¹
G = 0.007 S

K = κ / G = 0.014 / 0.007 = 2.0 cm⁻¹

Molar Conductance & Kohlrausch's Law - AJAX VI

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AJAX VI: Molar Conductance of Weak Electrolytes & Kohlrausch’s Law

1. Molar Conductance at a Given Concentration (Λ_m)

Definition: Conductance of all ions produced by one mole of an electrolyte in a given volume of solution.

Λ_m = κ × (1000 / M)

κ = specific conductance (S·cm⁻¹)
M = molarity (mol/L)

2. Molar Conductance at Infinite Dilution (Λ_m^∞)

At infinite dilution, electrolytes are completely dissociated. The conductance under this condition is denoted by:

Λ_m^∞ = limiting molar conductance

Note: For weak electrolytes like CH₃COOH, Λ_m^∞ is determined using **Kohlrausch’s Law**.

3. Kohlrausch’s Law

Statement:
*At infinite dilution, the molar conductivity of an electrolyte is equal to the sum of the individual ionic conductances of the cation and anion.*

Λ_m^∞ = λ^∞_cation + λ^∞_anion

Applications:

Calculate Λ_m^∞ of weak electrolytes using known strong electrolytes.
Determine degree of dissociation (α):
α = Λ_m / Λ_m^∞
Calculate dissociation constant (K_a):
K_a = (C × α²) / (1 − α)

4. Numerical Example

Q: Calculate Λ_m^∞ for CH₃COOH using:

Λ_m^∞(CH₃COONa) = 91 S·cm²·mol⁻¹
Λ_m^∞(HCl) = 426 S·cm²·mol⁻¹
Λ_m^∞(NaCl) = 126 S·cm²·mol⁻¹

Using Kohlrausch’s Law:

Λ_m^∞(CH₃COOH) = Λ_m^∞(CH₃COONa) + Λ_m^∞(HCl) − Λ_m^∞(NaCl)

= 91 + 426 − 126 = 391 S·cm²·mol⁻¹

AJAX VII - Faraday's Laws of Electrolysis

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AJAX VII – Faraday’s Laws of Electrolysis

1. Faraday’s First Law of Electrolysis

Statement: The mass \( m \) of a substance deposited or liberated at an electrode is directly proportional to the quantity of electricity \( Q \) passed.

Mathematical form:

\[
m = Z \cdot Q = Z \cdot I \cdot t
\]

\( m \) = mass of substance (g)
\( Z \) = electrochemical equivalent (g/C)
\( I \) = current in amperes (A)
\( t \) = time in seconds (s)

Example:
What mass of copper will be deposited by a current of 2 A in 30 minutes? \( Z_{\text{Cu}} = 0.000329 \, \text{g/C} \)
\( t = 30 \times 60 = 1800 \, \text{s} \)
\( m = Z \cdot I \cdot t = 0.000329 \times 2 \times 1800 = \boxed{1.1844 \, \text{g}} \)

2. Faraday’s Second Law of Electrolysis

Statement: When the same quantity of electricity is passed through different electrolytes, the masses of substances liberated are proportional to their chemical equivalents.

Mathematical form:

\[
\frac{m_1}{m_2} = \frac{E_1}{E_2}
\]

\( E \) = Equivalent weight = \( \frac{\text{Molar mass}}{n} \)

Example:
If 2.7 g of Al is deposited, how much Cu will be deposited by the same charge?
\( E_{\text{Al}} = 9, \quad E_{\text{Cu}} = 31.5 \)
\[
\frac{m_{\text{Cu}}}{2.7} = \frac{31.5}{9} \Rightarrow m_{\text{Cu}} = \frac{31.5}{9} \times 2.7 = \boxed{9.45 \, \text{g}}
\]

3. Relation: Faraday’s Constant, Avogadro's Number, and Charge on an Electron

The relationship is given by:

\[
F = N_A \cdot e
\]

\( F \) = Faraday's constant = \( 96500 \, \text{C/mol} \)
\( N_A \) = Avogadro's number = \( 6.022 \times 10^{23} \, \text{mol}^{-1} \)
\( e \) = Charge on one electron = \( 1.602 \times 10^{-19} \, \text{C} \)

Calculation:

\[
F = (6.022 \times 10^{23}) \times (1.602 \times 10^{-19}) \approx \boxed{96485 \, \text{C/mol}}
\]

AJAX VIII – Batteries and Fuel Cells

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AJAX – VIII: Batteries (Primary & Secondary Cells)

1. Primary Cell – Leclanché Cell

Type: Dry cell (non-rechargeable)
Structure: Zinc container (anode), carbon rod (cathode), MnO₂ and NH₄Cl paste as electrolyte
Electrode Reactions:

Anode: Zn → Zn²⁺ + 2e⁻
Cathode: 2MnO₂ + 2NH₄⁺ + 2e⁻ → Mn₂O₃ + 2NH₃ + H₂O

Use: Flashlights, clocks, remote controls

2. Primary Cell – Mercury Cell

Type: Button cell (non-rechargeable)
Structure: Zinc (anode), HgO (cathode), electrolyte: KOH + ZnO paste
Electrode Reactions:

Anode: Zn + 2OH⁻ → ZnO + H₂O + 2e⁻
Cathode: HgO + H₂O + 2e⁻ → Hg + 2OH⁻

Use: Hearing aids, watches, calculators

3. Secondary Cell – Lead Storage Battery

Type: Rechargeable
Structure: Lead (Pb) anode, PbO₂ cathode, sulfuric acid (H₂SO₄) electrolyte
Discharging Reactions:

Anode: Pb + SO₄²⁻ → PbSO₄ + 2e⁻
Cathode: PbO₂ + SO₄²⁻ + 4H⁺ + 2e⁻ → PbSO₄ + 2H₂O

Charging: Reverse of discharging reactions

Use: Inverters, automobiles, backup power

4. Fuel Cell – Hydrogen-Oxygen Fuel Cell

Type: Continuous fuel supply, eco-friendly
Structure: H₂ and O₂ gases, electrodes: porous carbon with catalysts, electrolyte: NaOH or KOH
Electrode Reactions:

Anode: H₂ + 2OH⁻ → 2H₂O + 2e⁻
Cathode: ½O₂ + H₂O + 2e⁻ → 2OH⁻

Overall Reaction:

H₂ + ½O₂ → H₂O (ΔG < 0)

Use: Spacecrafts, submarines, clean energy systems

AJAX IX – Corrosion Notes

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AJAX IX: Corrosion

1. Concept of Corrosion

Corrosion is the gradual destruction or deterioration of metals due to chemical or electrochemical reactions with their environment.

Example: Rusting of iron in presence of moisture and oxygen.

Chemical Representation:

Fe + O₂ + H₂O → Fe₂O₃·xH₂O (Rust)

2. Electrochemical Mechanism of Corrosion (Rusting of Iron)

In moist air, iron acts as an anode and undergoes oxidation, while oxygen acts as the cathode and is reduced.

At Anode (oxidation):

Fe → Fe²⁺ + 2e⁻

At Cathode (reduction):

O₂ + 4H⁺ + 4e⁻ → 2H₂O

The Fe²⁺ ions react with water and oxygen to form hydrated ferric oxide (rust).

Fe²⁺ + O₂ + H₂O → Fe₂O₃·xH₂O

3. Factors Affecting Corrosion

Presence of moisture: Increases corrosion rate
Presence of electrolytes: Salts accelerate corrosion
Impurities in metal: Create electrochemical cells
pH of environment: Acidic medium increases corrosion
Temperature: High temperatures accelerate corrosion

4. Prevention of Corrosion

Barrier protection: Painting, greasing, oiling
Galvanization: Coating iron with zinc
Alloying: Using corrosion-resistant alloys like stainless steel
Cathodic protection: Connecting the metal to a more reactive metal (sacrificial anode)
Electroplating: Coating the metal surface with a thin layer of another metal

3. Chemical Kinetics

Chemical Kinetics – AJAX Notes

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AJAX – Chemical Kinetics

1. Meaning of Chemical Kinetics

Chemical kinetics is the branch of chemistry that deals with the rate of chemical reactions, the mechanism by which they occur, and the factors influencing them.

Fast reactions: Occur almost instantaneously (e.g., precipitation, acid-base reactions).

Slow reactions: Take a longer time (e.g., rusting, radioactive decay).

2. Rate of a Reaction

Average Rate:

\[
\text{Average Rate} = \frac{\Delta [\text{Reactant or Product}]}{\Delta t}
\]

Instantaneous Rate:

\[
\text{Instantaneous Rate} = \lim_{\Delta t \to 0} \frac{\Delta [\text{Reactant}]}{\Delta t} = \frac{d[\text{R}]}{dt}
\]

Graphical Representation: Slope of concentration vs. time graph gives the rate at a point.

3. Factors Affecting Rate of Reaction

Surface Area: Greater surface area increases rate.
Nature of Reactants: Ionic reactions are faster than covalent ones.
Concentration: Higher concentration increases collision frequency.
Temperature: Rate increases with temperature due to more kinetic energy.
Catalyst: Lowers activation energy, increases rate.
Radiation: UV or sunlight can initiate or speed up reactions.

4. Order and Molecularity of a Reaction

Order: Sum of powers of concentration in the rate law.

\[
\text{Rate} = k[A]^x[B]^y \Rightarrow \text{Order} = x + y
\]

Molecularity: Number of reacting species involved in an elementary step (1, 2, or 3).

5. Rate Law and Specific Rate Constant

Rate Law: Expresses rate as a function of reactant concentrations.

Specific Rate Constant (k): The rate when all reactants have unit concentration.

6. Integrated Rate Equations and Half-Life

Zero Order Reaction:

\[
[R] = [R]_0 - kt
\]
\[
t_{1/2} = \frac{[R]_0}{2k}
\]

First Order Reaction:

\[
[R] = [R]_0 e^{-kt}
\quad \text{or} \quad
\ln[R] = \ln[R]_0 - kt
\]
\[
t_{1/2} = \frac{0.693}{k}
\]

7. Collision Theory (Elementary Idea)

Reactions occur due to collisions between reacting particles. For a reaction to occur:

Collisions must be effective (correct orientation and sufficient energy).
Only a fraction of collisions lead to reaction (those exceeding activation energy).

8. Threshold Energy and Activation Energy

Activation Energy (Ea): Minimum energy needed for a reaction to occur.

Threshold Energy: Total energy required by reactants to convert into products.

\[
\text{Threshold Energy} = \text{Activation Energy} + \text{Energy of Reactants}
\]

9. Arrhenius Equation

Relates rate constant \( k \) with temperature and activation energy:

\[
k = A e^{-E_a/RT}
\]
\[
\ln k = \ln A - \frac{E_a}{RT}
\]

\( A \) = frequency factor
\( E_a \) = activation energy
\( R \) = gas constant (8.314 J/mol·K)
\( T \) = temperature in Kelvin

AJAX-I – Introduction to Chemical Kinetics

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AJAX-I: Introduction to Chemical Kinetics

1. Meaning of Chemical Kinetics

Chemical kinetics is the branch of chemistry that deals with the study of the rate of chemical reactions, the factors affecting the rate, and the mechanism by which the reaction occurs.

Rate of reaction refers to the change in concentration of reactants or products with time:

\[
\text{Rate} = -\frac{d[\text{Reactant}]}{dt} = \frac{d[\text{Product}]}{dt}
\]

2. Scope and Importance of Chemical Kinetics

Helps in understanding the reaction mechanism.
Important in designing industrial chemical processes for maximum efficiency.
Essential in fields like pharmacology, biochemistry, and environmental science.
Enables prediction and control of explosive or very slow reactions.

3. Slow and Fast Reactions (Explained in Terms of Bonds)

Reactions are classified based on the time they take:

Fast reactions: Complete in fractions of a second (e.g., acid-base reactions, precipitation). Bonds break and form quickly.
Slow reactions: Take hours, days or years (e.g., rusting of iron). Involve stable bonds or multiple steps.

Explanation:

Reaction rate depends on the nature and strength of chemical bonds:

Weak or ionic bonds: Easier to break → faster reactions.
Strong covalent bonds: Harder to break → slower reactions.

Example:

Neutralization: HCl + NaOH → fast (ionic)
Rusting: Fe + O₂ + H₂O → slow (involves multiple covalent bond rearrangements)

AJAX-II – Rate of Reaction

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AJAX-II: Rate of Reaction

1. Definition of Rate of Reaction

The rate of a chemical reaction is the change in concentration of a reactant or product per unit time.

\[
\text{Rate} = -\frac{d[\text{Reactant}]}{dt} = \frac{d[\text{Product}]}{dt}
\]

The negative sign for reactants indicates a decrease in concentration over time.

2. Representation in Terms of Reactants and Products

For a general reaction:

\[
aA + bB \rightarrow cC + dD
\]

The rate is given by:

\[
\text{Rate} = -\frac{1}{a} \frac{d[A]}{dt} = -\frac{1}{b} \frac{d[B]}{dt} = \frac{1}{c} \frac{d[C]}{dt} = \frac{1}{d} \frac{d[D]}{dt}
\]

3. Determination of Rate Graphically

By plotting concentration vs. time:

Instantaneous Rate: Slope of the tangent at a specific time.
Average Rate: Slope of the chord between two points.

\[
\text{Average Rate} = \frac{\Delta[\text{R}]}{\Delta t}, \quad
\text{Instantaneous Rate} = \frac{d[\text{R}]}{dt}
\]

4. Instantaneous and Average Rate

Average Rate: Measured over a time interval \( \Delta t \).
Instantaneous Rate: Measured at a particular instant, using the derivative of concentration with respect to time.

5. Factors Affecting Rate of Reaction

Concentration: Higher concentration leads to faster reaction due to more collisions.
Temperature: Higher temperature increases kinetic energy and reaction rate.
Surface Area: Finer particles have more surface area, increasing the rate.
Nature of Reactants: Ionic reactions are faster than covalent ones.
Catalyst: Lowers activation energy and speeds up reaction.
Light (Radiation): Some reactions are photo-initiated (e.g., photosynthesis).

AJAX-III – Law of Mass Action

AJAX-III: Law of Mass Action

1. Statement of the Law of Mass Action

The Law of Mass Action states that:

"The rate of a chemical reaction is directly proportional to the product of the active masses (concentrations) of the reactants, each raised to a power equal to its stoichiometric coefficient in the balanced chemical equation."

2. Meaning of Active Mass

Active mass refers to the effective concentration of a substance that takes part in a chemical reaction.

For gases: active mass is taken as **partial pressure** or **mol/L**.
For solutions: active mass = molar concentration (mol/L).

3. General Form of the Law

For a general reaction:

\[
aA + bB \rightarrow cC + dD
\]

The rate of reaction is given by:

\[
\text{Rate} = k [A]^a [B]^b
\]

\( k \) = rate constant
\( [A], [B] \) = molar concentrations of reactants
\( a, b \) = stoichiometric coefficients of reactants

4. Example

Consider the reaction:

\[
2NO + O_2 \rightarrow 2NO_2
\]

According to the law of mass action:

\[
\text{Rate} \propto [NO]^2 [O_2]
\quad \Rightarrow \quad
\text{Rate} = k[NO]^2[O_2]
\]

AJAX-IV – Effect of Concentration on Reaction Rate

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AJAX-IV: Effect of Concentration on Rate of Reaction

1. Qualitative Explanation Based on Law of Mass Action

According to the Law of Mass Action, the rate of a chemical reaction increases as the concentration of the reactants increases.

This is because higher concentration leads to more frequent and effective collisions between reactant molecules.

2. Rate Law and General Rate Equation

The rate law is an experimentally determined expression that relates the rate of a reaction to the concentrations of its reactants.

\[
\text{Rate} = k [A]^n
\]

\( k \) = rate constant
\( [A] \) = concentration of reactant A
\( n \) = order of the reaction with respect to A

If more than one reactant is involved, the equation becomes:

\[
\text{Rate} = k [A]^x [B]^y
\]

\( x \), \( y \) = orders of reaction with respect to A and B respectively
Total order = \( x + y \)

3. Relationship Between Rate, Concentration, and Rate Constant

The rate of reaction is directly proportional to the product of the concentrations of reactants, each raised to their respective orders:

\[
\text{Rate} \propto [\text{Reactants}]^{\text{order}}
\]

At constant temperature:

\[
\frac{\text{Rate}_1}{\text{Rate}_2} = \left( \frac{[\text{A}]_1}{[\text{A}]_2} \right)^n
\]

This relation can be used to determine the order \( n \) experimentally by comparing rates at different concentrations.

4. Key Notes

\( k \) is independent of concentration, but depends on temperature.
\( n \) is determined experimentally, not necessarily equal to stoichiometric coefficients.
Unit of \( k \) depends on the order of the reaction.

AJAX-V – Order of a Reaction

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AJAX-V: Order of a Reaction

1. Meaning of Order of Reaction

The order of a chemical reaction is the sum of the powers of concentration terms in the experimentally determined rate law.

\[
\text{Rate} = k [A]^x [B]^y \quad \Rightarrow \quad \text{Order} = x + y
\]

2. Relation Between Order and Stoichiometric Coefficients

The order of a reaction is not always equal to the stoichiometric coefficients. It is determined experimentally.

Example:

\[
H_2 + I_2 \rightarrow 2HI, \quad \text{Rate} = k[H_2][I_2] \Rightarrow \text{Order} = 2
\]

3. Order as an Experimental Quantity

Order is determined using initial rate data. Compare rates at different concentrations to find the order \( n \):

\[
\frac{\text{Rate}_1}{\text{Rate}_2} = \left( \frac{[A]_1}{[A]_2} \right)^n
\]

4. Zero Order Reaction

Rate is independent of concentration.

\[
\text{Rate} = k, \quad [A] = [A]_0 - kt
\]
\[
\text{Unit of } k = \text{mol} \cdot \text{L}^{-1} \cdot \text{s}^{-1}
\]

5. First Order Reaction

Rate Law:

\[
\text{Rate} = k[A]
\]

Integrated Rate Equation:

Starting with:

\[
\frac{d[A]}{dt} = -k[A]
\]
\[
\Rightarrow \int_{[A]_0}^{[A]} \frac{d[A]}{[A]} = -k \int_0^t dt
\Rightarrow \ln\left(\frac{[A]_0}{[A]}\right) = kt
\]

Expression for Concentration:

\[
[A] = [A]_0 e^{-kt}
\]

Units of Rate Constant (k):

\[
k = \frac{1}{t} \Rightarrow \text{unit} = \text{s}^{-1}
\]

6. Half-life Period (\( t_{1/2} \))

Definition:

The time required for the concentration of a reactant to become half of its initial value.

Derivation (First Order):

\[
\ln\left(\frac{[A]_0}{[A]_0/2}\right) = kt_{1/2} \Rightarrow \ln(2) = kt_{1/2}
\Rightarrow t_{1/2} = \frac{0.693}{k}
\]

Note: For a first-order reaction, half-life is independent of initial concentration.

7. Characteristics of First Order Reactions

Rate depends only on the concentration of one reactant.
Graph of \(\ln[A]\) vs. time is linear (slope = -k).
\(t_{1/2}\) is constant.
Rate constant (k) is independent of initial concentration.

8. Solved Problem – First Order Rate Equation

Q: A first order reaction has a rate constant \( k = 1.386 \times 10^{-3} \, \text{s}^{-1} \). What is the half-life?

Solution:

\[
t_{1/2} = \frac{0.693}{k} = \frac{0.693}{1.386 \times 10^{-3}} = 500 \, \text{s}
\]

Q: The concentration of a reactant drops from 1.0 M to 0.25 M in 2 hours. Find rate constant.

Solution:

\[
\ln\left(\frac{[A]_0}{[A]}\right) = kt = \ln\left(\frac{1.0}{0.25}\right) = \ln(4) = 1.386
\]
\[
k = \frac{1.386}{2 \times 3600} = 1.925 \times 10^{-4} \, \text{s}^{-1}
\]

AJAX-VI – Molecularity of a Reaction

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AJAX-VI: Molecularity of a Reaction

1. Meaning of Molecularity

Molecularity of a reaction is the number of reactant molecules (atoms or ions) that collide simultaneously in a single step of a chemical reaction.

Physical picture: It gives an idea of how many species must come together at the molecular level for a reaction to occur.

It is always a whole number: 1, 2, or 3.
It is defined for elementary (one-step) reactions only.

2. Relation Between Molecularity, Order, and Rate

For a simple (elementary) reaction:

\[
A + B \rightarrow \text{Products}
\]
\[
\text{Rate} = k[A][B]
\]

Here:

Molecularity = 2 (bimolecular)
Order = 1 + 1 = 2

So, for elementary reactions:

\[
\text{Order} = \text{Molecularity}
\]

But for complex (multi-step) reactions, order may be different and must be determined experimentally.

3. Differences Between Order and Molecularity

Feature	Order	Molecularity
Definition	Sum of powers of concentration terms in rate law	Number of molecules taking part in the elementary step
Determination	Experimental	Theoretical (from mechanism)
Applicable To	Overall or complex reactions	Elementary steps only
Can it be fractional?	Yes (e.g., 1.5)	No – always a whole number
Can it be zero?	Yes (e.g., zero order reactions)	No – must be ≥ 1

4. Examples

Unimolecular reaction:

\[
\text{N}_2\text{O}_5 \rightarrow 2NO_2 + \frac{1}{2}O_2
\]
Molecularity = 1

Bimolecular reaction:

\[
\text{H}_2 + I_2 \rightarrow 2HI
\]
Molecularity = 2

AJAX-VII – The Concept of Energy

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AJAX-VII: The Concept of Energy in Chemical Reactions

1. Exothermic and Endothermic Reactions

Exothermic reaction: Releases energy (usually as heat); products have less energy than reactants.
Endothermic reaction: Absorbs energy; products have more energy than reactants.

\[
\Delta H = E_{\text{products}} - E_{\text{reactants}}
\]

\( \Delta H < 0 \): Exothermic
\( \Delta H > 0 \): Endothermic

2. Energy Barrier and Threshold Energy

Reactant molecules must possess a minimum energy to react – this is the threshold energy.

The difference between the threshold energy and the average energy of the reactants is called the activation energy (\(E_a\)).

\[
E_a = E_{\text{threshold}} - E_{\text{reactants}}
\]

3. Formation of Activated Complex

During the reaction, reactant molecules form an unstable intermediate with high energy called the activated complex (or transition state).

It exists momentarily at the peak of the energy barrier.
May break down to form products or revert to reactants.

4. Activation Energy

It is the minimum energy required by reactants to form the activated complex and initiate a chemical reaction.

Reactions with low activation energy are faster; those with high activation energy are slower.

5. Effect of Catalyst on Activation Energy

Catalysts provide an alternate reaction path with lower activation energy.
This increases the number of effective collisions, thus increasing the rate of reaction.

\[
E_{a,\text{catalysed}} < E_{a,\text{uncatalysed}} \]

6. Energy Profile Diagram Summary (Descriptive)

Uncatalyzed Reaction: Higher peak (more activation energy)

Catalyzed Reaction: Lower peak (less activation energy)

ΔH: Same for both since it depends only on initial and final states

AJAX-VIII – Collision Theory

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AJAX-VIII: Collision Theory

1. Conditions for a Chemical Change

Reactant particles must come into close contact.
They must collide with each other.
However, not all collisions lead to a reaction.

2. Effective Collisions

For a collision to result in a chemical reaction, it must be effective.

Conditions for effectiveness:

Sufficient energy: The particles must collide with energy ≥ activation energy \( E_a \).
Proper orientation: Molecules must align correctly for bond formation or breaking.

3. Energy Barrier and Activated Complex

Even in exothermic reactions, energy is needed to overcome the activation energy barrier to start the reaction.

This leads to the formation of an activated complex, a high-energy, unstable transition state.

\[
\text{Activation Energy (} E_a \text{)} = \text{Energy of Activated Complex} - \text{Energy of Reactants}
\]

4. Energy Profile Diagrams

Exothermic Reaction

• Reactants start at a higher energy level than products.
• Energy is released (ΔH < 0).
• Peak = activated complex. Difference between peak and reactants = \( E_a \).

Endothermic Reaction

• Products are at a higher energy level than reactants.
• Energy is absorbed (ΔH > 0).
• Still requires activation energy to reach the activated complex.

\[
\Delta H = E_{\text{products}} - E_{\text{reactants}}
\quad
E_a = E_{\text{activated complex}} - E_{\text{reactants}}
\]

5. Summary of Key Concepts

Collision Theory: Reactions occur when particles collide with enough energy and correct orientation.
Activated Complex: Temporary high-energy intermediate.
Activation Energy: Minimum energy required for a reaction to proceed.
Exothermic: \( \Delta H < 0 \), energy released.
Endothermic: \( \Delta H > 0 \), energy absorbed.

AJAX-IX – Mechanism of a Reaction

AJAX-IX: Mechanism of a Reaction

1. Meaning of Reaction Mechanism

A reaction mechanism is the step-by-step sequence of elementary reactions by which an overall chemical change occurs.

Elementary reaction: A single-step reaction that occurs in one molecular event.
Complex reaction: Composed of two or more elementary steps.
Overall reaction: The sum of all elementary steps (net reaction).

2. Rate-Determining Step (RDS)

In a multi-step reaction, the slowest elementary step is called the rate-determining step.

This step controls the overall rate of the reaction.

\[
\text{Overall Rate} = \text{Rate of the slowest step}
\]

3. Relation Between Rate Law and Mechanism

The rate law is based on the molecularity of the rate-determining step, not the overall reaction.

Example:

Reaction: \( 2NO + O_2 \rightarrow 2NO_2 \)

Mechanism:

Step 1 (slow): \( 2NO \rightarrow N_2O_2 \)
Step 2 (fast): \( N_2O_2 + O_2 \rightarrow 2NO_2 \)

Rate law from slow step:

\[
\text{Rate} = k[NO]^2
\]

Order: 2 (from rate-determining step)

4. Units of Rate Constant

The unit of rate constant \( k \) depends on the order of reaction.

Zero Order Reaction

\[
\text{Rate} = k \Rightarrow k = \text{mol·L}^{-1}\text{s}^{-1}
\]

First Order Reaction

\[
\text{Rate} = k[A] \Rightarrow k = \text{s}^{-1}
\]

Second Order Reaction

\[
\text{Rate} = k[A]^2 \Rightarrow k = \text{mol}^{-1}\text{L}\cdot\text{s}^{-1}
\]

5. Summary

Elementary reaction: One-step, simple reaction.
Complex reaction: Multiple steps, includes intermediates.
Mechanism: Stepwise representation of the reaction.
Rate-determining step: Slowest step; controls overall rate.
Rate law: Based on RDS, not entire reaction equation.
Order: Sum of powers of reactants in rate law.
Units of \( k \): Change with reaction order.

AJAX-X – Effect of Temperature: Arrhenius Equation

AJAX-X: Effect of Temperature on Rate Constant

1. Arrhenius Equation

The rate constant \( k \) increases with temperature, described by the Arrhenius equation:

\[
k = Ae^{-E_a / RT}
\]

Where:

\( k \) = rate constant
\( A \) = frequency factor or pre-exponential factor
\( E_a \) = activation energy (J/mol or kJ/mol)
\( R \) = universal gas constant (8.314 J/mol·K)
\( T \) = absolute temperature (in Kelvin)

2. Logarithmic Form

Taking natural logarithm on both sides:

\[
\ln k = \ln A - \frac{E_a}{RT}
\]

In base-10 logarithm:

\[
\log k = \log A - \frac{E_a}{2.303RT}
\]

3. Graphical Interpretation

Plotting \( \log k \) vs \( \frac{1}{T} \) gives a straight line with:

Slope = \( -\frac{E_a}{2.303R} \)
Intercept = \( \log A \)

\[
\text{Slope} = -\frac{E_a}{2.303R} \Rightarrow E_a = -\text{Slope} \times 2.303R
\]

4. Temperature and Collisions

Higher temperature → more molecules have energy ≥ \( E_a \).
This leads to a higher rate of effective collisions and faster reaction.
Rule of thumb: A 10°C rise roughly doubles the reaction rate.

5. Two-Temperature Arrhenius Equation

If rate constants \( k_1 \) and \( k_2 \) are known at two temperatures \( T_1 \) and \( T_2 \), we use:

\[
\log\left(\frac{k_2}{k_1}\right) = \frac{E_a}{2.303R} \left(\frac{1}{T_1} - \frac{1}{T_2}\right)
\]

6. Example Numerical (Setup)

Problem: At 300 K, rate constant \( k_1 = 1.5 \times 10^{-3} \). At 310 K, \( k_2 = 3.0 \times 10^{-3} \). Find \( E_a \).

Solution:

\[
\log\left(\frac{3.0 \times 10^{-3}}{1.5 \times 10^{-3}}\right) = \frac{E_a}{2.303 \times 8.314} \left(\frac{1}{300} - \frac{1}{310}\right)
\]

Solve for \( E_a \) after calculating both sides.

7. Summary Table

Equation: \( k = Ae^{-E_a/RT} \)
Log form: \( \log k = \log A - \frac{E_a}{2.303RT} \)
Slope of graph: \( -\frac{E_a}{2.303R} \)
Temperature ↑ → Rate ↑ due to more effective collisions

4. d- and f- Block Elements

d- and f-Block Elements

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d- and f-Block Elements

1. Position in the Periodic Table

- d-Block: Groups 3 to 12 (transition elements)
- f-Block: Lanthanides and Actinides, placed separately at the bottom

2. Occurrence and Electronic Configuration

General Configuration of d-block:

\[
(n-1)d^{1-10}ns^{0-2}
\]

- 3d series: Sc (Z = 21) to Zn (Z = 30)
- Elements are generally found in ores as oxides or sulfides

3. General Characteristics of 3d Transition Elements

Metallic character: Good conductors, malleable, ductile
Ionization enthalpy: Moderate and increases across the series
Oxidation states: Variable due to close energy of \( (n-1)d \) and \( ns \) orbitals
Ionic radii: Decrease slightly across the period (due to poor shielding)
Colour of ions: Due to d-d electronic transitions (incomplete d-orbitals)
Catalytic property: Used as catalysts (e.g., Fe in Haber’s process)
Magnetic properties: Due to unpaired electrons; paramagnetism shown
Interstitial compounds: Trap small atoms (H, C, N) in the lattice
Alloy formation: With other metals, forming strong and corrosion-resistant alloys

4. Potassium Dichromate (\( \text{K}_2\text{Cr}_2\text{O}_7 \))

Preparation: From chromite ore by fusion and acidification

Important Reaction:

\[
4FeCr_2O_4 + 8Na_2CO_3 + 7O_2 \rightarrow 8Na_2CrO_4 + 2Fe_2O_3 + 8CO_2
\]

Acidification: \( 2Na_2CrO_4 + H_2SO_4 \rightarrow \text{K}_2\text{Cr}_2\text{O}_7 + H_2O \)

Properties:

Strong oxidizing agent in acidic medium
Cr in +6 oxidation state

5. Potassium Permanganate (\( \text{KMnO}_4 \))

Preparation: By oxidation of \( \text{MnO}_2 \) with KOH and O₂ or KNO₃

\[
2MnO_2 + 4KOH + O_2 \rightarrow 2KMnO_4 + 2H_2O
\]

Properties:

Dark purple solid; Mn in +7 state
Strong oxidizing agent in all media (acidic, basic, neutral)
Used in water purification, titrations, and antiseptics

Lanthanoids, Actinoids and d-Block Elements

Lanthanoids, Actinoids, and d-Block Elements

1. d-Block Elements (3d, 4d, 5d Series)

Electronic configuration: General form:

\[
(n-1)d^{1-10}ns^{0-2}
\]

3d: Sc to Zn
4d: Y to Cd
5d: La/Hf to Hg

Key trends:

Metallic Character: High; all are good conductors
Atomic & Ionic Radii: Decrease slightly across a series
Ionization Enthalpy: Increases across series
Oxidation States: Variable due to similar energy levels of d and s orbitals
Variable Valency: Results from different d-electron participation
Coloured Compounds: Due to d-d transitions of unpaired d electrons
Complex Formation: Due to small size and high charge density of cations
Alloy Formation: With other transition and non-transition metals

2. Lanthanoids (4f-series)

Atomic Numbers: 57 (La) to 71 (Lu)

General configuration:

\[
[Xe]4f^{1-14}5d^{0-1}6s^2
\]

Properties:

Lanthanoid Contraction: Gradual decrease in atomic & ionic radii due to poor shielding by 4f electrons
Oxidation State: Mostly +3
Colour: Ions are coloured due to f-f transitions
Magnetism: Paramagnetic behavior due to unpaired 4f electrons

3. Actinoids (5f-series)

Atomic Numbers: 89 (Ac) to 103 (Lr)

General configuration:

\[
[Rn]5f^{1-14}6d^{0-1}7s^2
\]

Properties:

Radioactive: All actinoids are radioactive
Variable Oxidation States: Common states: +3, +4, +5, +6
Colour: Due to f-f transitions (similar to lanthanoids)
Complex Formation: High charge and small size → strong tendency to form complexes
Actinoid Contraction: Similar to lanthanoid contraction

4. Summary Table

Property	3d Series	4d Series	5d Series	Lanthanoids	Actinoids
Block	d	d	d	f	f
Common Oxidation State	+2 to +7	+2 to +8	+2 to +8	+3	+3, +4, +5, +6
Complex Formation	Yes	Yes	Yes	Moderate	High
Magnetism	Paramagnetic	Paramagnetic	Paramagnetic	Strong (unpaired 4f)	Very strong

f-Block Elements (4f & 5f Series)

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f-Block Elements – Lanthanoids and Actinoids

1. Electronic Configuration

Lanthanoids (4f series): Atomic numbers 57–71
Actinoids (5f series): Atomic numbers 89–103

General configuration:

\[
\text{Lanthanoids: } [Xe]4f^{1-14}5d^{0-1}6s^2 \\
\text{Actinoids: } [Rn]5f^{1-14}6d^{0-1}7s^2
\]

2. Atomic and Ionic Radii

Decrease gradually across the series (lanthanoid contraction)
Actinoids show a similar contraction called actinoid contraction

3. Oxidation States

Lanthanoids: Mostly +3
Actinoids: Variable oxidation states from +3 to +6 due to 5f orbital availability

4. Coloured Compounds

Due to f-f transitions of unpaired electrons
Less intense than d-d transitions (Laporte-forbidden)

5. Complex Formation

Actinoids form more stable and complex compounds than lanthanoids
Due to greater availability of 5f orbitals for bonding

6. Alloy Formation

Lanthanoids form alloys with metals like Fe, Al, Mg
Used in metallurgy and nuclear industries

7. Lanthanoid Contraction and Its Consequences

Definition: Steady decrease in ionic radii from Ce to Lu

Cause: Poor shielding of 4f electrons

Consequences:

Similarity in size of 4d and 5d elements (Zr and Hf)
Decrease in basicity of lanthanoid hydroxides
Separation becomes difficult due to similar ionic radii

8. Chemical Reactivity

Oxygen: Form oxides on heating
Hydrogen: Form hydrides
Halogens: Form trihalides like LnCl₃
Sulphur: Form sulphides (e.g., Ln₂S₃)
Nitrogen: Form nitrides (e.g., LnN)
Carbon: Form carbides
Water: React slowly to liberate hydrogen gas

9. Comparison: Lanthanoids vs Actinoids

Property	Lanthanoids (4f)	Actinoids (5f)
Electron Filling	4f orbitals	5f orbitals
Oxidation States	Mainly +3	+3 to +6
Radioactivity	Generally non-radioactive	All radioactive
Complex Formation	Less pronounced	More prominent
Shielding Effect	Poor (4f)	Very poor (5f)

Potassium Permanganate (KMnO₄)

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Potassium Permanganate (KMnO₄)

1. Structure and Shape

Chemical formula: KMnO₄
Shape of MnO₄⁻ ion: Tetrahedral
Oxidation state of Mn: +7

2. Extraction from Pyrolusite (MnO₂)

Step 1: Conversion of MnO₂ to Potassium Manganate (K₂MnO₄)

\[
2MnO_2 + 4KOH + O_2 \rightarrow 2K_2MnO_4 + 2H_2O
\]

Step 2: Oxidation of manganate to permanganate

\[
3K_2MnO_4 + 2CO_2 \rightarrow 2KMnO_4 + MnO_2 + 2K_2CO_3
\]

3. Oxidising Nature

In Acidic Medium: Mn⁷⁺ reduces to Mn²⁺

\[
MnO_4^- + 8H^+ + 5e^- \rightarrow Mn^{2+} + 4H_2O
\]

With FeSO₄:
\[
5Fe^{2+} + MnO_4^- + 8H^+ \rightarrow 5Fe^{3+} + Mn^{2+} + 4H_2O
\]
With Oxalic acid \((COOH)_2\cdot2H_2O\):
\[
5(COOH)_2 + 2MnO_4^- + 6H^+ \rightarrow 10CO_2 + 2Mn^{2+} + 8H_2O
\]
With KI:
\[
10I^- + 2MnO_4^- + 16H^+ \rightarrow 2Mn^{2+} + 5I_2 + 8H_2O
\]

In Basic Medium: MnO₄⁻ reduces to MnO₂

\[
MnO_4^- + 2H_2O + 3e^- \rightarrow MnO_2 + 4OH^-
\]

With KI:
\[
2MnO_4^- + H_2O + I^- \rightarrow 2MnO_2 + I_2 + 2OH^-
\]

In Neutral Medium: MnO₄⁻ still reduces to MnO₂

With H₂S:
\[
2MnO_4^- + 3H_2S \rightarrow 2MnO_2 + 3S + 2H_2O
\]

4. Use in Redox Titration

Acts as a self-indicator due to purple colour of MnO₄⁻ and colourless Mn²⁺ product
Used in:
- Titration of Fe²⁺ (ferrous)
- Oxalic acid titrations

Example: Redox titration of FeSO₄ with KMnO₄ in acid medium

\[
5Fe^{2+} + MnO_4^- + 8H^+ \rightarrow 5Fe^{3+} + Mn^{2+} + 4H_2O
\]

Potassium Dichromate (K₂Cr₂O₇)

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Potassium Dichromate (K₂Cr₂O₇)

1. Structure and Shape

Chemical formula: K₂Cr₂O₇
Dichromate ion (Cr₂O₇²⁻): Consists of two tetrahedral CrO₄ units sharing one oxygen
Chromium oxidation state: +6

2. Extraction from Chromite Ore (FeCr₂O₄)

Step 1: Fusion of chromite with Na₂CO₃ and O₂ to form sodium chromate

\[
4FeCr_2O_4 + 8Na_2CO_3 + 7O_2 \rightarrow 8Na_2CrO_4 + 2Fe_2O_3 + 8CO_2
\]

Step 2: Conversion of sodium chromate to sodium dichromate

\[
2Na_2CrO_4 + H_2SO_4 \rightarrow Na_2Cr_2O_7 + Na_2SO_4 + H_2O
\]

\[
Na_2Cr_2O_7 + 2KCl \rightarrow K_2Cr_2O_7 + 2NaCl
\]

3. Oxidising Nature

In Acidic Medium: Cr₂O₇²⁻ reduces to Cr³⁺

\[
Cr_2O_7^{2-} + 14H^+ + 6e^- \rightarrow 2Cr^{3+} + 7H_2O
\]

In Basic Medium: Converted to chromate ion (CrO₄²⁻) and acts less effectively

\[
Cr_2O_7^{2-} + 2OH^- \rightarrow 2CrO_4^{2-} + H_2O
\]

In Neutral Medium: Less reactive, requires catalyst or heat for oxidation

4. Use in Redox Titration

Strong oxidising agent in acidic medium
Used in redox titrations with Fe²⁺ and other reducing agents

Example:

\[
6Fe^{2+} + Cr_2O_7^{2-} + 14H^+ \rightarrow 6Fe^{3+} + 2Cr^{3+} + 7H_2O
\]

5. Interconversion of Chromate and Dichromate (Effect of pH)

In Acidic Medium:

\[
2CrO_4^{2-} + 2H^+ \rightleftharpoons Cr_2O_7^{2-} + H_2O
\]

In Basic Medium:

\[
Cr_2O_7^{2-} + 2OH^- \rightleftharpoons 2CrO_4^{2-} + H_2O
\]

Yellow colour: Chromate (\(CrO_4^{2-}\)) in basic medium
Orange colour: Dichromate (\(Cr_2O_7^{2-}\)) in acidic medium

5. Coordination Compounds

Coordination Compounds

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Coordination Compounds

1. Concepts and Definitions

Complex/Coordination Compound: Compound formed by a central metal ion surrounded by ligands.
Ligands: Ions or molecules that donate a lone pair to the central atom (e.g., NH₃, Cl⁻, CN⁻)
Coordination Number (CN): Number of ligand donor atoms bonded to the central atom.
Oxidation Number: Charge on the metal ion if all ligands are removed as neutral or anionic species.

2. IUPAC Nomenclature of Mononuclear Complexes

Cation first, then anion
Ligands are named alphabetically (not by charge)
Use prefixes: di-, tri-, tetra- for number of identical ligands
Examples:

\[
[Co(NH_3)_6]Cl_3 \rightarrow \text{Hexaamminecobalt(III) chloride}
\]
\[
K_2[PtCl_6] \rightarrow \text{Potassium hexachloroplatinate(IV)}
\]

3. Types of Isomerism

Structural Isomerism: Linkage, Ionisation, Coordination, Solvate
Stereoisomerism: Geometrical and Optical
Example: \([Co(NH_3)_4Cl_2]^+\) has cis- and trans- isomers

4. Bonding in Coordination Compounds

Werner's Theory

Primary valency: ionic (oxidation number)
Secondary valency: coordination number (directly bonded ligands)

Valence Bond Theory (VBT)

Explains geometry based on hybrid orbitals
sp³: tetrahedral, dsp²: square planar, d²sp³: octahedral

Crystal Field Theory (CFT)

Explains colour, magnetism based on d-orbital splitting
Octahedral field: \(d\)-orbitals split into \(t_{2g}\) and \(e_g\)

\[
\Delta_o = E(e_g) - E(t_{2g}) \quad \text{(Octahedral crystal field splitting)}
\]

5. Colour and Magnetic Properties

Colour arises from d-d electronic transitions
Magnetism depends on unpaired electrons
Paramagnetic: unpaired e⁻ | Diamagnetic: all electrons paired

6. Shapes of Complexes

Tetrahedral: [ZnCl₄]²⁻ (sp³)
Square Planar: [Ni(CN)₄]²⁻ (dsp²)
Octahedral: [Fe(CN)₆]³⁻ (d²sp³)

7. Importance of Coordination Compounds

Qualitative Analysis: Detection of metal ions (e.g., Ni²⁺ with DMG)
Extraction: Cyanide process for gold and silver extraction
Biological Systems: Hemoglobin (Fe), Chlorophyll (Mg), Vitamin B₁₂ (Co)

Coordination Compounds - Part I

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Coordination Compounds – Part I

1. Definition of Coordination Compounds

Coordination compounds (or complex compounds) are those in which a central metal atom or ion is bonded to a group of surrounding ligands through coordinate (dative) bonds.

\[
[Co(NH_3)_6]Cl_3 \quad \text{(Hexaamminecobalt(III) chloride)}
\]

2. Difference Between Double Salts and Coordination Compounds

Double salts: Dissociate completely in water to give all ions. E.g., Mohr's salt: \(FeSO_4 \cdot (NH_4)_2SO_4 \cdot 6H_2O\)
Coordination compounds: Do not dissociate into all constituent ions in solution.
Example: In \([Cu(NH_3)_4]SO_4\), the complex ion \([Cu(NH_3)_4]^{2+}\) stays intact.

3. Ligands

Ligands are atoms, ions, or molecules that donate a lone pair of electrons to the central metal ion.

Monodentate: Donates one pair (e.g., Cl⁻, NH₃)
Bidentate: Donates two pairs (e.g., \(C_2O_4^{2-}\), ethylenediamine)
Tridentate to Hexadentate: Multiple donor atoms
Polydentate: Ligands with more than one donor atom
Chelating ligands: Form ring-like structures with the metal ion

4. Coordination Number

Number of ligand atoms directly bonded to the central atom/ion.

Example: In \([Co(NH_3)_6]^{3+}\), the coordination number is 6.

5. Oxidation Number of Central Atom

It is calculated by treating all ligands as neutral or charged, summing their contributions, and solving for the metal's charge.

Example: In \([Fe(CN)_6]^{4-}\)

\[
x + 6(-1) = -4 \Rightarrow x = +2
\]

6. IUPAC Nomenclature Rules

Name cation first, then anion
Name ligands alphabetically before metal
Use prefixes (di-, tri-, tetra-, etc.)
Negative ligands end in "o" (e.g., chloro, cyano)
Neutral ligands use common names (e.g., NH₃ = ammine, H₂O = aqua)
Oxidation state of metal shown in Roman numerals

\[
[Pt(NH_3)_2Cl_2] \rightarrow \text{Diamminedichloroplatinum(II)}
\]
\[
K_3[Fe(CN)_6] \rightarrow \text{Potassium hexacyanoferrate(III)}
\]

Isomerism in Coordination Compounds

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Isomerism in Coordination Compounds

1. What is Isomerism?

Isomerism is the phenomenon where two or more compounds have the same chemical formula but different arrangements of atoms or bonds, resulting in different physical or chemical properties.

2. Types of Isomerism

A. Structural Isomerism

Occurs due to differences in the structure or connectivity of ligands.

Ionisation isomerism: Exchange of ligands inside and outside the coordination sphere.

\[
[Co(NH_3)_5Br]SO_4 \quad \text{vs} \quad [Co(NH_3)_5SO_4]Br
\]

Hydrate isomerism: Different numbers of water molecules inside or outside the coordination sphere.

\[
[Cr(H_2O)_6]Cl_3, \quad [Cr(H_2O)_5Cl]Cl_2 \cdot H_2O, \quad [Cr(H_2O)_4Cl_2]Cl \cdot 2H_2O
\]

Linkage isomerism: Same ligand bonds through different atoms.

Example: \(\text{NO}_2^-\) can bind via N (nitro) or O (nitrito).

\[
[Co(NH_3)_5(NO_2)]^{2+} \quad \text{vs} \quad [Co(NH_3)_5(ONO)]^{2+}
\]

Coordination isomerism: Exchange of ligands between cationic and anionic complexes.

\[
[Co(NH_3)_6][Cr(CN)_6] \quad \text{vs} \quad [Cr(NH_3)_6][Co(CN)_6]
\]

Ligand isomerism: Different isomers of the ligand itself (e.g., structural isomers).

B. Stereoisomerism

Occurs due to different spatial arrangements of ligands.

Geometrical isomerism: Ligands occupy different positions around the metal ion (cis–trans).

\[
[Pt(NH_3)_2Cl_2] \rightarrow \text{cis and trans forms}
\]

Optical isomerism: Non-superimposable mirror images (chiral compounds)

\[
[Co(en)_3]^{3+} \quad \text{(exists as d- and l- isomers)}
\]

3. Summary Table

Type	Subtypes	Example
Structural	Ionisation	\([Co(NH_3)_5Br]SO_4\)
Structural	Linkage	\([Co(NH_3)_5(NO_2)]^{2+}\) vs \([Co(NH_3)_5(ONO)]^{2+}\)
Stereoisomerism	Geometrical	\([Pt(NH_3)_2Cl_2]\)
Stereoisomerism	Optical	\([Co(en)_3]^{3+}\)

Valence Bond Theory – Coordination Compounds

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Valence Bond Theory (VBT) – Coordination Compounds

1. Introduction

Valence Bond Theory explains the bonding in coordination compounds based on the hybridization of orbitals in the central metal ion. It helps predict geometry and magnetic properties.

2. Types of Hybridization and Geometry

sp³: Tetrahedral geometry
dsp²: Square planar geometry
d²sp³: Inner orbital octahedral (low spin)
sp³d²: Outer orbital octahedral (high spin)

3. Inner vs Outer Orbital Complexes

Inner Orbital Complex (Low Spin)

Uses inner (n−1)d orbitals for hybridization (d²sp³).

Example: \([Fe(CN)_6]^{4-}\)

Fe²⁺: 3d⁶
CN⁻ is a strong field ligand → pairing occurs → low spin → d²sp³ → diamagnetic

Outer Orbital Complex (High Spin)

Uses outer nd orbitals for hybridization (sp³d²).

Example: \([FeF_6]^{3-}\)

Fe³⁺: 3d⁵
F⁻ is a weak field ligand → no pairing → high spin → sp³d² → paramagnetic

4. Magnetic Character Prediction

Based on number of unpaired electrons
Paramagnetic: At least one unpaired electron
Diamagnetic: All electrons paired
Magnetic moment: Calculated using formula:

\[
\mu = \sqrt{n(n+2)} \text{ BM}, \quad \text{where } n = \text{number of unpaired electrons}
\]

5. Geometrical Summary

Hybridization	Geometry	Example	Spin Type
d²sp³	Octahedral	[Fe(CN)₆]⁴⁻	Low Spin
sp³d²	Octahedral	[FeF₆]³⁻	High Spin
dsp²	Square Planar	[Ni(CN)₄]²⁻	Low Spin
sp³	Tetrahedral	[NiCl₄]²⁻	High Spin

Crystal Field Theory – Coordination Compounds

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Crystal Field Theory (CFT)

1. Introduction

Crystal Field Theory (CFT) explains the bonding and properties of coordination compounds by treating metal-ligand interactions as purely electrostatic.

2. Crystal Field Splitting in Octahedral Complexes

In an octahedral field, the five degenerate d-orbitals of the metal split into two sets:

t_2g (lower energy): d_xy, d_yz, d_zx
e_g (higher energy): d_z², d_x²−y²

\[
\Delta_0 = \text{Octahedral crystal field splitting energy}
\]

Examples:

Weak field ligand (e.g., F⁻): High spin → electrons remain unpaired
Strong field ligand (e.g., CN⁻): Low spin → electrons pair up in t_2g

3. Crystal Field Splitting in Tetrahedral Complexes

In a tetrahedral field, the d-orbitals split in the opposite pattern:

e (lower energy): d_z², d_x²−y²
t₂ (higher energy): d_xy, d_yz, d_zx

\[
\Delta_t = \frac{4}{9} \Delta_0
\]

Since \(\Delta_t\) is smaller than \(\Delta_0\), tetrahedral complexes are usually high spin (no pairing).

4. Colour of Complexes

Color arises from the d–d electronic transitions between t_2g and e_g levels.

Wavelength absorbed depends on \(\Delta\)
Complementary color is observed

Example: \([Ti(H_2O)_6]^{3+}\) appears violet due to absorption in the green-yellow region.

5. Magnetic Character

Depends on the number of unpaired electrons after splitting
High spin: More unpaired electrons → paramagnetic
Low spin: Fewer/no unpaired electrons → less paramagnetic or diamagnetic
Magnetic moment calculated using:

\[
\mu = \sqrt{n(n+2)} \text{ BM}, \quad \text{where } n = \text{unpaired electrons}
\]

6. Summary Table

Complex	Field Strength	Spin Type	Magnetism
[Fe(CN)₆]⁴⁻	Strong	Low Spin	Diamagnetic
[FeF₆]³⁻	Weak	High Spin	Paramagnetic
[NiCl₄]²⁻	Weak	High Spin (Tetrahedral)	Paramagnetic

Stability of Coordination Compounds

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Stability of Coordination Compounds

1. What is Stability?

The stability of a coordination compound refers to its ability to remain intact without decomposition. There are two types:

Thermodynamic stability: Related to the equilibrium constant of the complex formation.
Kinetic stability: Related to the rate at which the complex forms or decomposes.

2. Stability Constant (Formation Constant)

The equilibrium constant for the formation of a complex from a metal ion and ligands is called the stability constant, denoted as \( K_f \).

Example: For the complex \([ML_n]\):

\[
M^{n+} + nL \rightleftharpoons [ML_n]^{n+}
\]
\[
K_f = \frac{[ML_n^{n+}]}{[M^{n+}][L]^n}
\]

A higher \( K_f \) value indicates greater stability of the complex.

3. Stepwise Formation Constants

When ligands are added one at a time, we define stepwise stability constants \( K_1, K_2, \dots \)

\[
M + L \rightleftharpoons ML \quad K_1 = \frac{[ML]}{[M][L]}
\]
\[
ML + L \rightleftharpoons ML_2 \quad K_2 = \frac{[ML_2]}{[ML][L]}
\]
\[
\text{Overall: } \beta_n = K_1 \cdot K_2 \cdot \cdots \cdot K_n
\]

\( \beta_n \) is the overall stability constant of the complex.

4. Factors Affecting Stability

Charge on metal ion: Higher charge increases stability due to stronger electrostatic attraction.
Size of metal ion: Smaller ions form more stable complexes.
Basicity of ligands: More basic ligands donate electrons more easily → more stable complexes.
Chelate effect: Ligands with multiple donor atoms (like EDTA) form more stable complexes.
Nature of solvent: Polar solvents stabilize ionic complexes better.

5. Importance of Stability Constants

Helps predict if a reaction will proceed forward or not.
Used in quantitative analysis and separation of metal ions.
Essential in designing metal-ligand complexes in medicine and catalysis.

Importance and Uses of Coordination Compounds

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Importance and Uses of Coordination Compounds

1. Analytical Chemistry

Coordination compounds are widely used in qualitative and quantitative analysis:

Complexometric titrations: EDTA is commonly used to estimate metal ions.
Detection of metal ions: Formation of colored complexes helps in identifying specific metal ions.

2. Biological Importance

Many biological molecules are coordination compounds:

Hemoglobin: An iron coordination complex that transports oxygen in blood.
Chlorophyll: A magnesium coordination compound essential for photosynthesis.
Vitamin B₁₂: Contains a cobalt center important for enzymatic activity.

3. Industrial Applications

Catalysts: Coordination compounds are used as catalysts in various chemical reactions, e.g., in hydrogenation and polymerization.
Electroplating: Metal complexes are used to deposit metals like nickel and chromium on surfaces.
Preparation of pure metals: Complexes help in refining and extraction.

4. Medicinal Uses

Drugs: Cisplatin, a platinum complex, is used in cancer chemotherapy.
Diagnostic agents: Some coordination compounds are used in medical imaging techniques.

5. Environmental Uses

Removal of pollutants: Coordination compounds can bind heavy metals and help in water purification.
Soil analysis: Complexes help detect trace metals in soils.

6. Summary

Coordination compounds play vital roles in:

Qualitative and quantitative chemical analysis
Biological systems as essential molecules
Industrial catalysis and metallurgy
Medicine and environmental chemistry

6. Haloalkanes and Haloarenes.

Haloalkanes - Chemistry Notes

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Haloalkanes: Notes

1. General Formula

Haloalkanes have the general formula: \( C_nH_{2n+1}X \), where \( X \) is a halogen (F, Cl, Br, I).

2. Nomenclature

Common Name: Alkyl halides (e.g., methyl chloride)
IUPAC Name: Halogenoalkanes (e.g., chloromethane)
Prefix: Fluoro-, Chloro-, Bromo-, Iodo-

3. Classification

Primary (\(1^\circ\)) – Halogen attached to a C bonded with 1 other C.
Secondary (\(2^\circ\)) – Halogen on a C bonded with 2 other C atoms.
Tertiary (\(3^\circ\)) – Halogen on a C bonded with 3 other C atoms.

4. Nature of C–X Bond

Polar in nature due to electronegativity difference.
Bond strength: C–F > C–Cl > C–Br > C–I
Bond length increases from C–F to C–I.

5. Physical Properties

Boiling point increases with molecular weight and branching.
Insoluble in water, soluble in organic solvents.
Density: Haloalkanes are heavier than hydrocarbons.

6. Chemical Properties

Substitution Reactions

They undergo nucleophilic substitution:

\[
R{-}X + OH^- \rightarrow R{-}OH + X^-
\]

S_N1 Mechanism: Unimolecular, forms carbocation intermediate, rate = k[RX]
S_N2 Mechanism: Bimolecular, single-step, rate = k[RX][Nu⁻]

7. Optical Rotation

Chiral haloalkanes can rotate plane-polarized light.
Optical isomerism arises due to asymmetric carbon atom.
S_N2 reactions lead to inversion of configuration ("Walden inversion").

Haloarenes Notes

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Haloarenes

1. Basic Idea

Haloarenes are aromatic compounds in which one or more hydrogen atoms of the benzene ring are replaced by halogen atoms (F, Cl, Br, I).

General formula: \( C_6H_5X \), where \( X \) = halogen.

2. Nature of C–X Bond

The C–X bond in haloarenes is less reactive than in haloalkanes.
This is due to the partial double bond character from resonance:
\[
\text{C}_6\text{H}_5\text{–X} \leftrightarrow \text{C}_6\text{H}_4^{\delta-}\text{=X}^{\delta+}
\]
The bond length is shorter than a typical C–X single bond due to this character.

3. Substitution Reactions

Haloarenes undergo electrophilic substitution rather than nucleophilic substitution.

The halogen atom is deactivating but ortho-para directing due to resonance.

Common reactions: nitration, sulphonation, halogenation, Friedel–Crafts alkylation/acylation.
Example:
\[
C_6H_5Cl + HNO_3 \xrightarrow{H_2SO_4} o\text{-}ClC_6H_4NO_2 + p\text{-}ClC_6H_4NO_2
\]

Halogen Compounds: Uses & Environmental Effects

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Uses & Environmental Effects of Some Halogen Compounds

1. Dichloromethane (DCM) — \( CH_2Cl_2 \)

Uses: Solvent in paint removers, propellant in aerosols, degreasing agent.
Effects: Can cause dizziness, nausea; contributes to air pollution.

2. Trichloromethane (Chloroform) — \( CHCl_3 \)

Uses: Solvent, anesthetic (historically).
Effects: Harmful to liver and kidneys; possible carcinogen; contributes to ozone layer depletion.

3. Tetrachloromethane (Carbon Tetrachloride) — \( CCl_4 \)

Uses: Cleaning agent, fire extinguisher (old use).
Effects: Toxic to liver; depletes ozone; not biodegradable.

4. Iodoform — \( CHI_3 \)

Uses: Antiseptic in medical dressings.
Effects: Slightly toxic; less environmental harm but smells strongly.

5. Freons — e.g., \( CCl_2F_2 \)

Uses: Refrigerants, aerosol propellants.
Effects: Major cause of ozone depletion.

6. DDT (Dichlorodiphenyltrichloroethane)

Uses: Insecticide against mosquitoes and agricultural pests.
Effects: Bioaccumulates in food chain, toxic to wildlife and humans, banned in many countries.

Nature of C–X Bond

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Nature of the C–X Bond

1. Bond Characteristics

The C–X bond in haloalkanes is a **polar covalent bond**.
Halogen (X) is more electronegative than carbon:
\[
\text{C–X} \Rightarrow \delta^+ \text{C} - \delta^- \text{X}
\]
This polarity causes the carbon atom to become slightly electron-deficient, making it reactive toward nucleophiles.

2. Bond Strength & Length

Bond strength decreases as the size of halogen increases:
\[
\text{C–F} > \text{C–Cl} > \text{C–Br} > \text{C–I}
\]
Bond length increases down the group due to larger atomic radii of halogens.

3. Reactivity

Polar nature makes C–X bond a site for nucleophilic substitution.
Reactivity depends on bond enthalpy and inductive effects.

Naming Halogen Derivatives of Alkanes

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Naming Halogen Derivatives of Alkanes

1. Common System

Named as alkyl halides.
Example:
- \( CH_3Cl \): Methyl chloride
- \( CH_3CH_2Br \): Ethyl bromide
Prefix for number of halogens: di-, tri-, tetra-, etc.

2. IUPAC System

Halogens are treated as substituents on the parent alkane chain.
Use prefixes: fluoro, chloro, bromo, iodo.
Number the carbon chain to give halogens the lowest possible locants.
Example:
\[
CH_3CHClCH_3 \Rightarrow 2\text{-chloropropane}
\]

3. Di- and Tri-Halo Derivatives

Multiple identical halogen atoms: use di-, tri- with locants.
Example:
\[
CH_2ClCHClCH_3 \Rightarrow 1,2\text{-dichloropropane}
\]
Mixed halogen atoms: list alphabetically.
\[
CBrClFCH_3 \Rightarrow 1\text{-bromo-1-chloro-1-fluoroethane}
\]

Preparation of Haloalkanes

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Preparation of Haloalkanes

1. From Alkanes and Halogens

Haloalkanes can be prepared by free radical halogenation of alkanes in the presence of UV light or heat:

\[
CH_4 + Cl_2 \xrightarrow{hv} CH_3Cl + HCl
\]

Note: This method is not selective and leads to mixtures of products.

2. From Alkenes and Hydrogen Halides

Addition of hydrogen halide (HX) to an alkene:

\[
CH_2=CH_2 + HBr \rightarrow CH_3CH_2Br
\]

Follows Markovnikov's rule:
\[
CH_3CH=CH_2 + HCl \rightarrow CH_3CHClCH_3
\]

3. From Alcohols

Alcohols react with halogenating agents:

With \( PCl_5 \):
\[
CH_3CH_2OH + PCl_5 \rightarrow CH_3CH_2Cl + POCl_3 + HCl
\]
With \( PX_3 \) (e.g. \( PBr_3 \)):
\[
3CH_3OH + PBr_3 \rightarrow 3CH_3Br + H_3PO_3
\]
With \( SOCl_2 \) (best method, by retention):
\[
CH_3CH_2OH + SOCl_2 \rightarrow CH_3CH_2Cl + SO_2 + HCl
\]

4. Halide Exchange Method

Replacement of one halogen by another using halide salts:

Finkelstein Reaction

Alkyl chloride/bromide with NaI in acetone:

\[
CH_3CH_2Br + NaI \xrightarrow{acetone} CH_3CH_2I + NaBr
\]

Swarts Reaction

Used to prepare alkyl fluorides:

\[
RBr + AgF \rightarrow RF + AgBr
\]

5. From Silver Salt of Fatty Acids (Hunsdiecker Reaction)

Reaction of silver carboxylates with halogens gives alkyl halides:

\[
C_nH_{2n+1}COOAg + Br_2 \rightarrow C_nH_{2n+1}Br + CO_2 + AgBr
\]

Example:

\[
C_2H_5COOAg + Br_2 \rightarrow C_2H_5Br + CO_2 + AgBr
\]

Haloalkanes: Physical & Chemical Properties

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Haloalkanes: Physical and Chemical Properties

1. Physical Properties

State: Lower members (e.g. methyl halides) are gases; others are liquids or solids.
Melting and Boiling Points:
- Increase with molecular mass.
- Increase with stronger van der Waals forces.
- For isomers: straight chains & more polarizable halogens → higher b.p.
Solubility: Insoluble in water (non-polar); soluble in organic solvents like benzene and ether.

2. Chemical Properties

(a) Nucleophilic Substitution Reactions

Replacement of halogen by a nucleophile (\( Nu^- \)):

SN1 Mechanism (Unimolecular):

Two steps: formation of carbocation → attack by nucleophile.
Favored by tertiary halides.
Example:
\[
(CH_3)_3CBr \xrightarrow{H_2O} (CH_3)_3COH + HBr
\]

SN2 Mechanism (Bimolecular):

One-step, backside attack.
Favored by primary halides.
Example:
\[
CH_3CH_2Br + OH^- \rightarrow CH_3CH_2OH + Br^-
\]

(b) Reactions with Specific Reagents

Sodium hydroxide (aqueous):
\[
R-X + NaOH_{(aq)} \rightarrow R-OH + NaX
\]
Water (hydrolysis):
\[
R-X + H_2O \rightarrow R-OH + HX
\]
Sodium iodide (Finkelstein reaction):
\[
R-Cl + NaI \xrightarrow{acetone} R-I + NaCl
\]
Ammonia (alcoholic):
\[
R-X + NH_3 \rightarrow R-NH_2 + HX
\]
Primary amine (\( R'NH_2 \)):
\[
R-X + R'NH_2 \rightarrow R-NHR' + HX
\]
Secondary amine (\( R'_2NH \)):
\[
R-X + R'_2NH \rightarrow R-NR'_2 + HX
\]
Potassium cyanide (KCN):
\[
R-X + KCN \rightarrow R-CN + KX
\]
Silver cyanide (AgCN):
\[
R-X + AgCN \rightarrow R-NC + AgX
\]
(Forms isocyanide due to bonding via N)
Potassium nitrite (KNO₂):
\[
R-X + KNO_2 \rightarrow R-ONO + KX
\]
(Nitrite formation)
Silver nitrite (AgNO₂):
\[
R-X + AgNO_2 \rightarrow R-NO_2 + AgX
\]
(Nitroalkane formation via N-bonding)
Silver salt of fatty acid (Hunsdiecker):
\[
RCOOAg + Br_2 \rightarrow R-Br + CO_2 + AgBr
\]
Lithium aluminium hydride (LiAlH₄):
\[
R-X + 2[H] \xrightarrow{LiAlH_4} R-H + HX
\]
(Reduction to alkane)

Elimination & Metal Reactions of Haloalkanes

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Elimination and Metal Reactions of Haloalkanes

1. β-Elimination (Dehydrohalogenation)

When a haloalkane is heated with alcoholic KOH, it undergoes β-elimination to form an alkene.

\[
R-CH_2CHX-R' + alc.\ KOH \xrightarrow{\Delta} R-CH=CH-R' + KX + H_2O
\]

Saytzeff’s Rule

In elimination, the more substituted alkene is the major product.

\[
CH_3CHBrCH_3 + alc.\ KOH \rightarrow CH_2=CHCH_3 + CH_3CH=CH_2
\]
Major product: 2-butene (more substituted)

2. Reaction with Sodium – Wurtz Reaction

Haloalkanes react with sodium metal in dry ether to form higher alkanes.

\[
2R-X + 2Na \xrightarrow{dry\ ether} R-R + 2NaX
\]

Example:

\[
2CH_3Br + 2Na \rightarrow C_2H_6 + 2NaBr
\]

3. Reaction with Magnesium – Grignard Reagent Preparation

Haloalkanes react with magnesium in dry ether to form Grignard reagents.

\[
R-X + Mg \xrightarrow{dry\ ether} R-MgX
\]

Example:

\[
CH_3Br + Mg \xrightarrow{dry\ ether} CH_3MgBr
\]

Grignard reagents are highly reactive and used to synthesize alcohols, carboxylic acids, etc.

Chloroform and Iodoform

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Chloroform and Iodoform

1. Chloroform (CHCl₃)

Preparation

From Ethanol or Acetone with bleaching powder (\( Ca(OCl)_2 \)):

\[
CH_3CH_2OH + 4[O] \rightarrow CH_3COOH + 2H_2O
\]
\[
CH_3COCH_3 + 3Cl_2 + 4NaOH \rightarrow CHCl_3 + CH_3COONa + 3NaCl + 2H_2O
\]

Properties

Heavy, colorless, volatile liquid with sweet odor.
Immiscible with water, soluble in alcohol and ether.
Under sunlight, oxidizes to toxic phosgene:
\[
2CHCl_3 + O_2 \xrightarrow{hv} 2COCl_2 + 2HCl
\]
Stored in dark bottles with ethanol.
Used as anesthetic (now replaced due to toxicity).

2. Iodoform (CHI₃)

Preparation

Iodoform test: heating ethanol or acetone with iodine and NaOH gives yellow precipitate of iodoform.

\[
CH_3CH_2OH + I_2 + 4NaOH \rightarrow CHI_3 + HCOONa + 3NaI + 2H_2O
\]

Properties

Yellow crystalline solid with antiseptic smell.
Slightly soluble in water; more in ethanol and ether.
Antiseptic (due to slow release of iodine).
Test for methyl ketones or ethanol.

Freons & Preparation of Haloarenes

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Structure of Freons & Preparation of Haloarenes

1. Structure of Freons

Freons are chlorofluorocarbons (CFCs) used as refrigerants and propellants.
General formula: \( CCl_mF_nH_{4−m−n} \)
Common example: **Freon-12** = **CCl₂F₂** (Dichlorodifluoromethane)
Structure:
\[
F−C(F)(Cl)−Cl \quad \text{(tetrahedral geometry)}
\]
Stable, non-toxic, non-flammable (but harmful to ozone layer).

2. Preparation of Haloarenes

a) Sandmeyer’s Reaction

Aryl diazonium salts react with Cu(I) halide to give haloarenes.

\[
C_6H_5N_2^+Cl^- + CuCl \rightarrow C_6H_5Cl + N_2 \uparrow
\]
\[
C_6H_5N_2^+Cl^- + CuBr \rightarrow C_6H_5Br + N_2 \uparrow
\]

b) Gattermann’s Reaction

Similar to Sandmeyer’s but uses Cu powder and halogen (Cl₂ or Br₂) directly:

\[
C_6H_5N_2^+Cl^- + Cu + Cl_2 \rightarrow C_6H_5Cl + N_2 + CuCl
\]

c) Electrophilic Substitution

Halogenation of benzene in presence of Lewis acid (FeCl₃/AlCl₃):

\[
C_6H_6 + Cl_2 \xrightarrow{FeCl_3} C_6H_5Cl + HCl
\]
\[
C_6H_6 + Br_2 \xrightarrow{FeBr_3} C_6H_5Br + HBr
\]

Mechanism: electrophilic attack of \( Cl^+ \) or \( Br^+ \) on aromatic ring.
Reaction follows typical EAS pathway: generation of electrophile → π-complex → substitution.

Physical Properties of Haloalkanes and Haloarenes

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Physical Properties of Haloalkanes and Haloarenes

1. Physical State

Lower haloalkanes (e.g., methyl chloride, ethyl bromide) are gases at room temperature.
Higher haloalkanes and all haloarenes are usually liquids or solids.
Physical state depends on molecular weight and type of halogen.

2. Melting and Boiling Points

Boiling point increases with:
- Increasing molecular mass
- More polarizable halogens (I > Br > Cl > F)
Straight-chain isomers have higher b.p. than branched ones.
For isomeric haloarenes:
- Para-isomers have higher melting points (symmetry leads to better crystal packing)

Example (boiling point trend):

\[
CH_3Cl < CH_3Br < CH_3I \]

3. Solubility

Haloalkanes and haloarenes are:
- Insoluble in water (they cannot form hydrogen bonds)
- Soluble in non-polar solvents like benzene, ether, and CCl₄
As halogen increases, polarity increases, but water solubility still remains low.

Reason for water insolubility:

\[
\text{Energy released in forming new dipole-dipole interactions is less than energy needed to break hydrogen bonds in water}
\]

Haloarenes: Chemical Properties

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Haloarenes: Chemical Properties

1. Electrophilic Substitution Reactions

Directive Influence: Halogens are ortho/para-directing due to resonance, but deactivating due to –I effect.

a) Chlorination

\[
C_6H_5Cl + Cl_2 \xrightarrow{FeCl_3} o\text{-}ClC_6H_4Cl + p\text{-}ClC_6H_4Cl
\]

b) Nitration

\[
C_6H_5Cl + HNO_3 \xrightarrow{H_2SO_4} o\text{-}NO_2C_6H_4Cl + p\text{-}NO_2C_6H_4Cl
\]

c) Sulphonation

\[
C_6H_5Cl + H_2SO_4 \xrightarrow{heat} o\text{-}HSO_3C_6H_4Cl + p\text{-}HSO_3C_6H_4Cl
\]

General Mechanism (EAS)

Step 1: Generation of electrophile (e.g. \( Cl^+ \), \( NO_2^+ \))
Step 2: Attack on aromatic ring → formation of arenium ion
Step 3: Loss of proton restores aromaticity

2. Nucleophilic Substitution

Haloarenes are less reactive toward nucleophilic substitution due to delocalization of lone pairs on halogen.

a) Replacement by –OH

\[
C_6H_5Cl + NaOH \xrightarrow{300^\circ C,\ 300\ atm} C_6H_5OH + NaCl
\]

b) Replacement by –NH₂

\[
C_6H_5Cl + NH_3 \xrightarrow{Cu_2O,\ 200^\circ C} C_6H_5NH_2 + HCl
\]

Mechanism (Benzyne Intermediate)

Step 1: Elimination of halide → benzyne formation
Step 2: Nucleophilic attack on benzyne → substituted product

3. Reduction to Benzene

Haloarenes can be reduced using hydrogen and catalyst:

\[
C_6H_5Cl + 2[H] \xrightarrow{H_2,\ Pd} C_6H_6 + HCl
\]

4. Wurtz-Fittig and Fittig Reactions

a) Wurtz-Fittig Reaction

Reaction of haloarene and haloalkane with sodium in dry ether:

\[
C_6H_5Br + CH_3Br + 2Na \xrightarrow{dry\ ether} C_6H_5CH_3 + 2NaBr
\]

b) Fittig Reaction

Two haloarenes + sodium in dry ether → biphenyl:

\[
2C_6H_5Br + 2Na \xrightarrow{dry\ ether} C_6H_5{-}C_6H_5 + 2NaBr
\]

5. Reaction with Magnesium (Grignard Reagent Formation)

Haloarenes react with magnesium in dry ether to form Grignard reagents:

\[
C_6H_5Br + Mg \xrightarrow{dry\ ether} C_6H_5MgBr
\]

Grignard reagents are useful for forming C–C bonds.

6. Structure of DDT

DDT = Dichlorodiphenyltrichloroethane

Structural formula:
\[
(ClC_6H_4)_2CHCCl_3
\]

Two p-chlorophenyl groups and one trichloromethyl group.
Highly stable, non-biodegradable insecticide (now restricted due to environmental concerns).

7. Alcohols, Phenols and Ethers

Alcohols: Classification, Properties, and Uses

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Alcohols: Classification, Properties, and Uses

1. Classification of Alcohols

Monohydric alcohols: Contain one –OH group
General formula: \( C_nH_{2n+1}OH \)
Dihydric alcohols: Contain two –OH groups
General formula: \( C_nH_{2n}(OH)_2 \)
Polyhydric alcohols: Contain more than two –OH groups
Example: Glycerol, \( C_3H_5(OH)_3 \)

2. Structure and General Formula

Alcohols are derivatives of alkanes with one or more hydrogen atoms replaced by hydroxyl groups (–OH).

General formula for monohydric alcohols: \( R{-}OH \), where \( R \) is an alkyl group.

Example: Ethanol – \( \mathrm{CH_3CH_2OH} \)

3. Nomenclature

Common system: Alkyl group + "alcohol" (e.g., methyl alcohol)
IUPAC system: Replace "-e" in alkane name with "-ol" (e.g., methanol, ethanol)

4. Methods of Preparation of Primary Alcohols

From alkenes by hydration:
\[
\mathrm{CH_2=CH_2 + H_2O \xrightarrow{H_2SO_4} CH_3CH_2OH}
\]
From Grignard reagent:
\[
\mathrm{CH_3MgBr + H_2O \rightarrow CH_4 + Mg(OH)Br}
\]
(general method to prepare alcohols from alkyl halides)
Hydrolysis of alkyl halides:
\[
\mathrm{CH_3CH_2Cl + NaOH \rightarrow CH_3CH_2OH + NaCl}
\]
Reduction of carbonyl compounds: e.g., reduction of aldehydes by \( \mathrm{LiAlH_4} \)
From primary amines: via Hofmann rearrangement (outline only)
Industrial methods:
- Methanol by Bosch process: \( \mathrm{CO + 2H_2 \xrightarrow{catalyst} CH_3OH} \)
- Ethanol by fermentation of carbohydrates:
  \[
  \mathrm{C_6H_{12}O_6 \xrightarrow{yeast} 2 C_2H_5OH + 2 CO_2}
  \]

5. Physical Properties of Primary Alcohols

Generally liquids at room temperature
Boiling points are higher than corresponding alkanes due to hydrogen bonding
Soluble in water due to –OH group; solubility decreases with increasing alkyl chain length

6. Chemical Properties of Primary Alcohols

Combustion:
\[
\mathrm{C_2H_5OH + 3O_2 \rightarrow 2CO_2 + 3H_2O}
\]
Oxidation:
\[
\mathrm{CH_3CH_2OH \xrightarrow{[O]} CH_3CHO \xrightarrow{[O]} CH_3COOH}
\]
Reaction with sodium:
\[
2 \mathrm{C_2H_5OH} + 2 \mathrm{Na} \rightarrow 2 \mathrm{C_2H_5ONa} + H_2
\]
Esterification with acids: e.g.
\[
\mathrm{CH_3CH_2OH + CH_3COOH \xrightarrow{H^+} CH_3COOCH_2CH_3 + H_2O}
\]

7. Identification of Primary, Secondary, and Tertiary Alcohols

Using Lucas Test:

Tertiary alcohols: Immediate turbidity with Lucas reagent
Secondary alcohols: Turbidity after 5–10 minutes
Primary alcohols: No turbidity or after long time

Examples:
Primary: \( \mathrm{CH_3CH_2OH} \)
Secondary: \( \mathrm{CH_3CHOHCH_3} \)
Tertiary: \( \mathrm{(CH_3)_3COH} \)

8. Mechanism of Dehydration of Primary Alcohol

Dehydration occurs in the presence of concentrated sulfuric acid producing an alkene:

\[
\mathrm{C_2H_5OH} \xrightarrow[\text{conc. } H_2SO_4]{\Delta} \mathrm{C_2H_4} + H_2O
\]

Stepwise mechanism:

Protonation of –OH to form better leaving group (\( H_2O \))
Loss of water forming carbocation intermediate
Removal of proton from adjacent carbon → formation of double bond (alkene)

9. Uses of Alcohols (Methanol and Ethanol)

Methanol: Used as solvent, antifreeze, fuel, and raw material for formaldehyde production.
Ethanol: Used as alcoholic beverage, solvent, fuel additive (bioethanol), and antiseptic.

Alcohols - Classification and Properties

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Alcohols: Classification and Properties

1. Classification

Monohydric alcohols: Contain one –OH group
General formula: \( C_nH_{2n+1}OH \)
Dihydric alcohols: Contain two –OH groups
General formula: \( C_nH_{2n}(OH)_2 \)
Polyhydric alcohols: Contain more than two –OH groups
Example: Glycerol \( C_3H_5(OH)_3 \)

2. Structure and General Formula

Alcohols are derivatives of alkanes where a hydrogen atom is replaced by a hydroxyl group \( -OH \).

General formula (monohydric): \( R{-}OH \), where \( R \) is an alkyl group.

3. Nomenclature

Common system: Alkyl group + "alcohol"
e.g. Methyl alcohol
IUPAC system: Replace “-e” in alkane with “-ol”
e.g. Methanol \( CH_3OH \), Ethanol \( C_2H_5OH \)

4. Classification: Primary, Secondary, Tertiary Alcohols

Primary (1°): –OH group on carbon bonded to 1 other carbon
Example: \( CH_3CH_2OH \) (ethanol)
Secondary (2°): –OH group on carbon bonded to 2 other carbons
Example: \( CH_3CHOHCH_3 \) (isopropanol)
Tertiary (3°): –OH group on carbon bonded to 3 other carbons
Example: \( (CH_3)_3COH \) (tert-butanol)

5. Physical and Chemical Properties

Physical Properties

Hydrogen bonding → High boiling points
Soluble in water (lower alcohols)
Insoluble as chain length increases

Chemical Properties

Combustion: \( C_2H_5OH + 3O_2 \rightarrow 2CO_2 + 3H_2O \)
Oxidation:
- Primary: → Aldehyde → Acid
- Secondary: → Ketone
- Tertiary: → No simple oxidation

4.1 Classification: Primary, Secondary, Tertiary Alcohols

Based on the number of alkyl groups attached to the carbon bearing the –OH group:

Primary (1°) Alcohol

–OH group attached to a carbon bonded to one other carbon.

Example: Ethanol \( CH_3CH_2OH \)

CH₃−CH₂−OH

Secondary (2°) Alcohol

–OH group on a carbon bonded to two other carbon atoms.

Example: Isopropanol \( CH_3CHOHCH_3 \)

CH₃−CH(OH)−CH₃

Tertiary (3°) Alcohol

–OH group on a carbon bonded to three carbon atoms.

Example: Tert-butanol \( (CH_3)_3COH \)

(CH₃)₃C−OH

Alcohols - Methods of Preparation

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Alcohols: Methods of Preparation

1. Hydration of Alkenes

a) Direct Hydration

\[
CH_2=CH_2 + H_2O \xrightarrow{H_2SO_4} CH_3CH_2OH
\]
(Ethene to ethanol using acid catalyst)

b) Indirect Hydration (Oxymercuration-demercuration)

\[
CH_2=CH_2 \xrightarrow[NaBH_4]{Hg(OAc)_2,\ H_2O} CH_3CH_2OH
\]

c) Hydroboration-Oxidation

\[
CH_2=CH_2 \xrightarrow[B_2H_6]{THF} CH_3CH_2BH_2 \xrightarrow[OH^-]{H_2O_2} CH_3CH_2OH
\]

2. From Grignard’s Reagent

\[
R-MgX + H_2C=O \xrightarrow{ether} RCH_2OMgX \xrightarrow{H_2O} RCH_2OH
\]
(Grignard + formaldehyde gives primary alcohol)

CH₃MgBr + HCHO → CH₃CH₂OH

3. Hydrolysis of Alkyl Halides

\[
R-X + OH^- \rightarrow R-OH + X^-
\]
Example:
\[
CH_3CH_2Br + NaOH \rightarrow CH_3CH_2OH + NaBr
\]

4. Reduction of Carbonyl Compounds

a) Aldehyde to Primary Alcohol

\[
RCHO + 2[H] \xrightarrow{LiAlH_4} RCH_2OH
\]

b) Ketone to Secondary Alcohol

\[
RCOR' + 2[H] \rightarrow RCH(OH)R'
\]

5. From Primary Amines

By diazotization and hydrolysis:

\[
RNH_2 + HNO_2 \rightarrow ROH + N_2 + H_2O
\]

6. Manufacture of Methanol

Bosch Process:

\[
CO + 2H_2 \xrightarrow{ZnO/Cr_2O_3,\ 300^\circ C,\ 300\ atm} CH_3OH
\]

7. Manufacture of Ethanol (Fermentation)

From Glucose using Yeast:

\[
C_6H_{12}O_6 \xrightarrow{yeast} 2C_2H_5OH + 2CO_2
\]

Alcohols - Chemical Properties

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Properties of Alcohols

1. Acidic Nature of Alcohols

Alcohols can donate a proton (H⁺) from the hydroxyl group, showing weak acidity:

\[
ROH + Na \rightarrow RONa + \frac{1}{2}H_2 \uparrow
\]

2. Reaction with Sodium

Alcohols react with active metals like sodium to form alkoxides:

\[
2CH_3CH_2OH + 2Na \rightarrow 2CH_3CH_2ONa + H_2 \uparrow
\]

3. Esterification (with mechanism)

Alcohols react with carboxylic acids in the presence of an acid catalyst (H₂SO₄) to form esters.

\[
RCOOH + R'OH \xrightarrow{H^+} RCOOR' + H_2O
\]

Mechanism:

Protonation of carbonyl oxygen
Nucleophilic attack by alcohol
Proton transfers and loss of water
Formation of ester

4. Reaction with Hydrogen Halides (HX)

Alcohols react with hydrogen halides to form alkyl halides:

\[
R-OH + HX \rightarrow R-X + H_2O
\]
Example:
\[
CH_3CH_2OH + HBr \rightarrow CH_3CH_2Br + H_2O
\]

5. Reaction with PCl₃, PCl₅, and SOCl₂

\( 3ROH + PCl_3 \rightarrow 3RCl + H_3PO_3 \)
\( ROH + PCl_5 \rightarrow RCl + POCl_3 + HCl \)
\( ROH + SOCl_2 \rightarrow RCl + SO_2 + HCl \)

6. Reaction with Acid Chlorides and Acid Anhydrides

Alcohols react with acid chlorides to give esters:

\[
RCOCl + R'OH \rightarrow RCOOR' + HCl
\]

With acid anhydrides:

\[
(RCO)_2O + R'OH \rightarrow RCOOR' + RCOOH
\]

7. Oxidation

Primary alcohol → Aldehyde → Carboxylic acid
Secondary alcohol → Ketone
Tertiary alcohol → No oxidation (under mild conditions)

\[
CH_3CH_2OH \xrightarrow{[O]} CH_3CHO \xrightarrow{[O]} CH_3COOH
\]

8. Dehydration (with mechanism)

Alcohols dehydrate in acidic medium to give alkenes.

\[
CH_3CH_2OH \xrightarrow{H_2SO_4,\ 170^\circ C} CH_2=CH_2 + H_2O
\]

Mechanism:

Protonation of –OH group → forms a good leaving group (H₂O)
Loss of water forms carbocation
Elimination of proton from β-carbon → alkene

Uses of Alcohols

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Uses of Alcohols

Methanol (CH₃OH)

CH₃–OH

Used as a solvent and antifreeze.
Raw material for manufacturing formaldehyde.
Fuel in race cars and as biodiesel additive.

Ethanol (C₂H₅OH)

CH₃–CH₂–OH

Used in alcoholic beverages.
Solvent in pharmaceuticals and cosmetics.
Fuel additive (ethyl alcohol in petrol).
Antiseptic and disinfectant.

Glycerol (C₃H₅(OH)₃)

HO–CH₂–CH(OH)–CH₂–OH

Used in cosmetics and pharmaceutical creams.
Humectant in food and tobacco industries.
Raw material for nitroglycerin (explosives).

General Uses of Alcohols

Used as solvents for dyes, oils, resins, and gums.
Intermediate in chemical synthesis (esters, ethers, etc.).
Used as antifreeze agents in cooling and heating systems.
Used in the manufacture of perfumes and flavorings.
Fuel and fuel additives for internal combustion engines.
Antiseptics and disinfectants in medical applications.

Conversion of One Alcohol into Another

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Conversion of One Alcohol into Another

1. Primary Alcohol <=> Aldehyde <=> Carboxylic Acid

CH₃CH₂OH (Ethanol)

Oxidation: Primary alcohol is oxidized to aldehyde, then further to carboxylic acid.

\[
\mathrm{CH_3CH_2OH} \xrightarrow[\text{mild}]{[O]} \mathrm{CH_3CHO} \xrightarrow[\text{strong}]{[O]} \mathrm{CH_3COOH}
\]

CH₃CHO (Acetaldehyde)

CH₃COOH (Acetic Acid)

Reduction: Carboxylic acid or aldehyde reduced back to primary alcohol.

\[
\mathrm{CH_3COOH} \xrightarrow{[H]} \mathrm{CH_3CH_2OH}
\]

2. Primary Alcohol <=> Secondary Alcohol

CH₃CH₂OH (Ethanol)

Oxidation of primary alcohol to aldehyde, then rearrangement to secondary alcohol via Grignard reagent.

Example: Using Grignard reagent with formaldehyde

\[
\mathrm{CH_3MgBr} + \mathrm{HCHO} \xrightarrow{H^+} \mathrm{CH_3CH_2OH}
\]

3. Secondary Alcohol <=> Tertiary Alcohol

CH₃CHOHCH₃ (Isopropanol)

(CH₃)₃COH (Tert-butanol)

Secondary alcohol can be converted to tertiary alcohol by reaction with Grignard reagent:

\[
\mathrm{CH_3CHOHCH_3} + \mathrm{CH_3MgBr} \xrightarrow{H^+} (CH_3)_3COH
\]

4. Conversion via Halides

Alcohols can be converted to alkyl halides, which then can be converted to other alcohols by substitution.

\[
R{-}OH \xrightarrow{PCl_5} R{-}Cl \xrightarrow{NaOH} R{-}OH
\]

Lucas’ Test - Distinguishing Alcohols

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Lucas’ Test: Distinguishing Alcohols

Introduction

Lucas’ Test uses Lucas reagent — a mixture of concentrated hydrochloric acid and zinc chloride (\( \mathrm{ZnCl_2} \)) — to differentiate primary (1°), secondary (2°), and tertiary (3°) alcohols based on their reactivity.

Reaction Principle

Alcohol reacts with Lucas reagent to form alkyl chloride (\( R{-}Cl \)) via nucleophilic substitution:

\[
R{-}OH + \mathrm{HCl} \xrightarrow[\mathrm{ZnCl_2}]{\text{Lucas reagent}} R{-}Cl + H_2O
\]

The rate of reaction depends on the alcohol type:

Tertiary alcohols react immediately (cloudiness forms within seconds)
Secondary alcohols react within minutes (cloudiness forms in 5–10 min)
Primary alcohols react very slowly or not at room temperature (cloudiness after hours or no reaction)

Structures of Alcohols

Primary alcohol (1°): CH₃CH₂OH

Secondary alcohol (2°): CH₃CHOHCH₃

Tertiary alcohol (3°): (CH₃)₃COH

Mechanism (Simplified)

The reaction proceeds by the formation of a carbocation intermediate. The stability of the carbocation increases from primary to tertiary, explaining the reaction speed:

Tertiary carbocation: Most stable → fast reaction
Secondary carbocation: Moderately stable → moderate reaction time
Primary carbocation: Least stable → slow or no reaction

Summary Table

Alcohol Type	Structure Example	Reaction Time with Lucas Reagent	Observation
Primary (1°)	CH₃CH₂OH	Hours or no reaction	No immediate turbidity
Secondary (2°)	CH₃CHOHCH₃	5–10 minutes	Gradual turbidity
Tertiary (3°)	(CH₃)₃COH	Immediate	Instant turbidity (cloudiness)

Phenols: Classification, Properties, and Uses

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Phenols: Classification, Properties, and Uses

1. Classification and Nomenclature

Monohydric phenols: Contain one –OH group attached to aromatic ring.
Example: Phenol \( \mathrm{C_6H_5OH} \)
Dihydric phenols (Catechols): Contain two –OH groups.
Example: Catechol (1,2-dihydroxybenzene)
Polyhydric phenols: More than two –OH groups.
Example: Pyrogallol (1,2,3-trihydroxybenzene)

Nomenclature: The –OH group attached to benzene ring forms the suffix phenol or prefix hydroxy- if another principal group exists.

2. Methods of Preparation

From chlorobenzene (Fusion with aqueous NaOH):
\[
\mathrm{C_6H_5Cl + NaOH \xrightarrow{heat} C_6H_5OH + NaCl}
\]
From cumene (Cumene process):
\[
\mathrm{(CH_3)_2CHC_6H_5 \xrightarrow{O_2} cumene\ hydroperoxide \xrightarrow{acid} phenol + acetone}
\]
From diazonium salts:
\[
\mathrm{C_6H_5N_2^+Cl^- + H_2O \rightarrow C_6H_5OH + N_2 + HCl}
\]
From sodium phenoxide:
\[
\mathrm{C_6H_5ONa + HCl \rightarrow C_6H_5OH + NaCl}
\]

3. Physical Properties

Colorless crystalline solid with characteristic odor
Melting point: 40.5°C, boiling point: 181.7°C
Soluble in water due to hydrogen bonding, also soluble in organic solvents
Has weakly acidic nature compared to alcohols

4. Chemical Properties

Acidic nature: Phenol is more acidic than alcohols due to resonance stabilization of phenoxide ion.
\[
\mathrm{C_6H_5OH \rightleftharpoons C_6H_5O^- + H^+}
\]
Reaction with sodium hydroxide:
\[
\mathrm{C_6H_5OH + NaOH \rightarrow C_6H_5ONa + H_2O}
\]
Reaction with zinc dust: Phenol reduces zinc to form zinc phenoxide:
\[
\mathrm{2 C_6H_5OH + Zn \rightarrow (C_6H_5O)_2Zn + H_2}
\]
Electrophilic substitution reactions (occurs easily due to activation by –OH group):
- Nitration:
  \[
  \mathrm{C_6H_5OH + HNO_3 \rightarrow 2\text{-}nitrophenol + 4\text{-}nitrophenol}
  \]
- Halogenation:
  \[
  \mathrm{C_6H_5OH + 3Cl_2 \rightarrow 2,4,6\text{-}trichlorophenol + 3HCl}
  \]
- Sulphonation:
  \[
  \mathrm{C_6H_5OH + H_2SO_4 \rightarrow C_6H_4OH\text{-}SO_3H + H_2O}
  \]

5. Uses of Phenols

Used as antiseptics and disinfectants
Used in manufacture of plastics like Bakelite
Starting material for synthesis of pharmaceuticals and dyes
Used as antioxidants and preservatives

Preparation of Phenol

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Preparation of Phenol

1. From Diazonium Salt

Phenol is prepared by the hydrolysis of benzene diazonium chloride:

\( \mathrm{C_6H_5N_2^+Cl^- + H_2O \xrightarrow{warm} C_6H_5OH + N_2 + HCl} \)

Note: The diazonium salt is prepared by treating aniline with nitrous acid at 0–5°C.

2. From Chlorobenzene (Dow’s Process)

Chlorobenzene is fused with sodium hydroxide at high temperature and pressure to give sodium phenoxide, which on acidification yields phenol:

\( \mathrm{C_6H_5Cl + 2NaOH \xrightarrow{350^\circ C, 300 atm} C_6H_5ONa + NaCl + H_2O} \)

\( \mathrm{C_6H_5ONa + HCl \rightarrow C_6H_5OH + NaCl} \)

3. From Benzene Sulphonic Acid

Benzene sulphonic acid on fusion with sodium hydroxide produces sodium phenoxide, which upon acidification gives phenol:

\( \mathrm{C_6H_5SO_3H + 2NaOH \xrightarrow{fusion} C_6H_5ONa + Na_2SO_3 + H_2O} \)

\( \mathrm{C_6H_5ONa + HCl \rightarrow C_6H_5OH + NaCl} \)

Manufacture of Phenol from Cumene

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Manufacture of Phenol from Cumene (Cumene Process)

Overview

Phenol is industrially manufactured by the cumene process, which involves the following steps:

Alkylation of benzene with propylene to form cumene (isopropylbenzene).
Oxidation of cumene to cumene hydroperoxide.
Acid-catalyzed cleavage of cumene hydroperoxide to give phenol and acetone.

Step 1: Formation of Cumene

Benzene reacts with propylene in presence of an acid catalyst (like AlCl₃) to form cumene:

\( \mathrm{C_6H_6 + CH_2=CHCH_3 \xrightarrow{AlCl_3} C_6H_5CH(CH_3)_2} \)

Cumene: \( \mathrm{C_6H_5{-}CH(CH_3)_2} \) (Isopropylbenzene)

Step 2: Oxidation of Cumene

Cumene is oxidized by air to form cumene hydroperoxide:

\( \mathrm{C_6H_5CH(CH_3)_2 + O_2 \rightarrow C_6H_5C(CH_3)_2OOH} \)

Cumene hydroperoxide: \( \mathrm{C_6H_5{-}C(CH_3)_2{-}OOH} \)

Step 3: Acid-Catalyzed Cleavage

Cumene hydroperoxide is cleaved in acidic medium to give phenol and acetone:

\( \mathrm{C_6H_5C(CH_3)_2OOH \xrightarrow{H^+} C_6H_5OH + (CH_3)_2CO} \)

Phenol: \( \mathrm{C_6H_5OH} \)
Acetone: \( \mathrm{(CH_3)_2CO} \)

Summary Reaction

\( \mathrm{C_6H_6 + CH_2=CHCH_3 + \frac{1}{2} O_2 \xrightarrow{} C_6H_5OH + (CH_3)_2CO} \)

Phenol - Physical Properties

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Phenol: Physical Properties

1. Physical State

Phenol is a white crystalline solid at room temperature with a pleasant odor.

Phenol structural formula: \( \mathrm{C_6H_5OH} \)

2. Solubility

Phenol is moderately soluble in water due to hydrogen bonding between the hydroxyl group and water molecules.
Solubility in water decreases as temperature increases.
Phenol is soluble in organic solvents like alcohol, ether, and chloroform.

The hydroxyl group (\( -OH \)) in phenol forms hydrogen bonds, which contributes to its solubility in polar solvents.

Chemical Properties of Phenol

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Chemical Properties of Phenol

1. Acidic Character of Phenol

Phenol is weakly acidic due to the ability of the hydroxyl group (\( -OH \)) to release \( H^+ \) ions.
The acidic nature is enhanced by resonance stabilization of the phenoxide ion:

\( \mathrm{C_6H_5OH \rightleftharpoons C_6H_5O^- + H^+} \)

2. Reaction with Sodium Hydroxide

Phenol reacts with sodium hydroxide to form sodium phenoxide and water:

\( \mathrm{C_6H_5OH + NaOH \rightarrow C_6H_5ONa + H_2O} \)

3. Reaction with Sodium Metal

Phenol reacts with sodium metal to give sodium phenoxide and hydrogen gas:

\( \mathrm{2C_6H_5OH + 2Na \rightarrow 2C_6H_5ONa + H_2 \uparrow} \)

4. Reaction with Zinc

When heated with zinc dust, phenol forms benzene:

\( \mathrm{C_6H_5OH + Zn \xrightarrow{\Delta} C_6H_6 + ZnO} \)

5. Reaction with Acetyl Chloride and Acetic Anhydride

Phenol reacts with acetyl chloride or acetic anhydride to form esters (acetates):

\( \mathrm{C_6H_5OH + CH_3COCl \rightarrow C_6H_5OCOCH_3 + HCl} \)

\( \mathrm{C_6H_5OH + (CH_3CO)_2O \rightarrow C_6H_5OCOCH_3 + CH_3COOH} \)

6. Reaction with Phosphorus Pentachloride (PCl₅)

Phenol reacts with PCl₅ to form phenyl chloride and phosphorus oxychloride:

\( \mathrm{C_6H_5OH + PCl_5 \rightarrow C_6H_5Cl + POCl_3 + HCl} \)

7. Electrophilic Substitution Reactions

Phenol undergoes electrophilic substitution reactions readily due to activation by the –OH group:

Bromination: Forms 2,4,6-tribromophenol (white precipitate):

\( \mathrm{C_6H_5OH + 3Br_2 \rightarrow C_6H_2Br_3OH + 3HBr} \)

Nitration: Forms ortho- and para-nitrophenol with dilute nitric acid:

\( \mathrm{C_6H_5OH + HNO_3 \rightarrow o\text{-} and\ p\text{-}NO_2C_6H_4OH + H_2O} \)

Sulphonation: With concentrated sulfuric acid forms phenol sulfonic acid:

\( \mathrm{C_6H_5OH + H_2SO_4 \rightarrow C_6H_4SO_3H + H_2O} \)

8. Kolbe’s Reaction (Formation of Salicylic Acid)

Phenol reacts with carbon dioxide in presence of sodium hydroxide under pressure and heat to form sodium salicylate, which on acidification gives salicylic acid:

\( \mathrm{C_6H_5ONa + CO_2 \xrightarrow{P, \Delta} C_6H_4(OH)COONa} \)

\( \mathrm{C_6H_4(OH)COONa + HCl \rightarrow C_6H_4(OH)COOH + NaCl} \)

9. Reimer–Tiemann Reaction

Phenol reacts with chloroform in the presence of sodium hydroxide to form salicylaldehyde (o-hydroxybenzaldehyde):

\( \mathrm{C_6H_5OH + CHCl_3 + 3NaOH \rightarrow C_6H_4(OH)CHO + 3NaCl + 2H_2O} \)

10. Tests for Phenol

FeCl₃ Test: Phenol gives violet color with neutral ferric chloride due to complex formation.
Azo Dye Test: Phenol reacts with diazonium salts to form colored azo dyes.

Aliphatic Ethers - Notes

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Aliphatic Ethers: Notes

1. General Formula and Structure

Aliphatic ethers have the general formula: \( R{-}O{-}R' \), where \( R \) and \( R' \) are alkyl groups.

Example: Diethyl ether — \( CH_3CH_2{-}O{-}CH_2CH_3 \)

Structure consists of an oxygen atom bonded to two alkyl groups.

2. Nomenclature

Common system: Name the alkyl groups attached to oxygen in alphabetical order followed by "ether".
Example: Ethyl methyl ether.
IUPAC system: Name the two alkyl groups as substituents attached to oxygen, followed by "ether".
Example: Ethoxyethane (for diethyl ether).

3. Methods of Preparation

Dehydration of Alcohols: Primary alcohols on heating with acid catalyst (like sulfuric acid) form ethers:
\( 2R{-}OH \xrightarrow{H_2SO_4, \Delta} R{-}O{-}R + H_2O \)
Williamson Ether Synthesis: Reaction of alkoxide ion with primary alkyl halide:
\( R{-}O^- + R'-X \rightarrow R{-}O{-}R' + X^- \)
From Halides and Silver Oxide: Alkyl halides react with \( Ag_2O \) to form ethers.

4. Physical Properties

Usually colorless liquids with pleasant odors.
Lower boiling points than alcohols of similar molecular weight due to absence of hydrogen bonding.
Insoluble in water or moderately soluble depending on alkyl chain length.
Less dense than water.

5. Chemical Properties

Relatively inert to bases and dilute acids.
Can be cleaved by strong acids like HI and HBr forming alkyl halides:
\( R{-}O{-}R' + HI \rightarrow R{-}I + R'{-}OH \) (for asymmetrical ethers)
Burn with a pale, sooty flame.
Undergo substitution reaction on alkyl halide formation.

6. Uses

Diethyl ether used as a solvent and formerly as a general anesthetic.
Used as solvents for fats, oils, waxes, perfumes, etc.
Used as starting materials for the synthesis of various organic compounds.

Ethers: Structure, Preparation, and Properties

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Ethers: Structure, Preparation, and Properties

1. Structure of Ethers (Ethereal Group)

Ethers contain an oxygen atom bonded to two alkyl or aryl groups represented as \( R{-}O{-}R' \). The oxygen is sp³ hybridized with two lone pairs.

Example: Diethyl ether — \( CH_3CH_2{-}O{-}CH_2CH_3 \)

2. Preparation of Ethers

Williamson Ether Synthesis

Reaction of alkoxide ion with primary alkyl halide:

\( R{-}O^- + R'-X \rightarrow R{-}O{-}R' + X^- \)

Where \( R{-}O^- \) is an alkoxide ion and \( R'-X \) is a primary alkyl halide.

3. Physical Properties

Usually colorless, volatile liquids with characteristic ether odor.
Low boiling points compared to alcohols of similar molecular weight (due to lack of hydrogen bonding).
Generally immiscible with water, but lower ethers show some solubility.
Less dense than water.

4. Chemical Properties

Reaction with Chlorine: Ethers can form alkyl halides and chlorinated products under photochemical conditions.
Oxidation (Peroxide Formation): Ethers slowly form peroxides on exposure to air and light:
\( R{-}O{-}R + O_2 \xrightarrow{light} R{-}O{-}O{-}R \) (Peroxide)
Reaction with Hydrogen Iodide (HI): Cleavage of ethers into alkyl iodides and alcohols:
\( R{-}O{-}R' + 2 HI \rightarrow R{-}I + R'{-}I + H_2O \)
Reaction with Phosphorus Pentachloride (PCl₅): Converts ethers to alkyl chlorides:
\( R{-}O{-}R' + PCl_5 \rightarrow R{-}Cl + R'{-}Cl + POCl_3 \)

Aryl Ethers: Notes

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Aryl Ethers: Notes

1. Physical Properties

Typically colorless liquids or solids.
Lower boiling points compared to corresponding alcohols.
Generally insoluble or sparingly soluble in water due to non-polar aromatic ring.
More soluble in organic solvents like benzene, ether, etc.

2. Chemical Properties

Preparation of Anisole (Phenyl Methyl Ether)

Prepared by Williamson Ether Synthesis:

\( C_6H_5O^- + CH_3I \rightarrow C_6H_5OCH_3 + I^- \)

(Phenoxide ion reacts with methyl iodide)

Electrophilic Substitution Reactions

The ether oxygen activates the aromatic ring toward electrophilic substitution by donating electron density via resonance.

Halogenation: Substitution of hydrogen by halogen (e.g. chlorine) occurs mainly at the ortho and para positions.
Nitration: Reaction with nitrating mixture \( HNO_3/H_2SO_4 \) introduces a nitro group mainly at ortho and para positions.
Friedel-Crafts Alkylation/Acylation: Anisole undergoes Friedel-Crafts reactions favoring ortho and para positions.

3. Uses of Ethers

Used as solvents in organic reactions due to low reactivity (e.g. diethyl ether).
Diethyl ether historically used as an anesthetic.
Anisole and other aryl ethers serve as intermediates in perfumes and pharmaceuticals.
Used as starting materials for synthesis of various organic compounds.

8. Aldehydes, Ketones and Carboxylic Acids

Aldehydes and Ketones - Notes

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Aldehydes and Ketones

1. Nomenclature

Aldehydes: IUPAC name ends in "-al". Common names based on corresponding acids.
Example: \( CH_3CHO \) → Ethanal (acetaldehyde)
Ketones: IUPAC name ends in "-one".
Example: \( CH_3COCH_3 \) → Propanone (acetone)

2. Structure

Both aldehydes and ketones contain the carbonyl group \( >C=O \), where the carbon is sp² hybridized and planar.

Aldehyde: R–CHO
Ketone: R–CO–R'

3. Methods of Preparation

From alcohols:
- Primary alcohol → Aldehyde
- Secondary alcohol → Ketone
From acid chlorides:
\( RCOCl + H_2/Pd \rightarrow RCHO \) (Rosenmund reduction)
From nitriles:
\( RCN + SnCl_2/HCl \rightarrow RCHO \)
From alkenes:
Ozonolysis of alkenes produces aldehydes and ketones.

4. Physical Properties

Polar compounds with higher boiling points than alkanes but lower than alcohols.
Lower members are soluble in water due to hydrogen bonding with water molecules.
Have sharp pungent odours (especially aldehydes).

5. Chemical Properties

Nucleophilic Addition Reactions

Hydrogen cyanide (HCN): Forms cyanohydrins
Sodium bisulfite (NaHSO₃): Forms bisulfite adducts
Alcohols: Forms hemiacetals and acetals
Grignard reagent: Forms alcohols

Mechanism of Nucleophilic Addition

Nucleophile attacks the electrophilic carbon of the carbonyl group, pushing electrons onto oxygen.

R–C(=O)–R' + Nu⁻ → R–C(OH)(Nu)–R'

Oxidation

Aldehydes oxidize to carboxylic acids.
Ketones resist oxidation (only under strong conditions).

Reduction

Aldehydes → Primary alcohols
Ketones → Secondary alcohols

6. Reactivity of Alpha Hydrogen

The alpha hydrogen (adjacent to the carbonyl group) is acidic due to resonance stabilization of the enolate ion:

\( R{-}CH_2{-}C(=O){-}R' \rightleftharpoons R{-}CH^{-}{-}C(=O){-}R' \rightarrow \) Enolate

This property enables aldol condensation and related reactions.

7. Uses of Aldehydes and Ketones

Formaldehyde used in polymer and resin production.
Acetaldehyde used in perfumes and dyes.
Acetone (a ketone) used as industrial solvent and nail polish remover.
Used as intermediates in pharmaceutical and plastic industries.

Preparation of Aldehydes and Ketones

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Preparation of Aldehydes and Ketones

1. From Alcohols

- Primary alcohols are oxidized to aldehydes.
- Secondary alcohols are oxidized to ketones.

\( CH_3CH_2OH \xrightarrow{[O]} CH_3CHO + H_2O \) (Ethanol to Ethanal)
\( CH_3CHOHCH_3 \xrightarrow{[O]} CH_3COCH_3 \) (Isopropanol to Acetone)

2. From Alkenes (Ozonolysis)

Cleavage of double bonds by ozone followed by hydrolysis gives aldehydes/ketones.

\( RCH=CHR' \xrightarrow{O_3} RCHO + R'CHO \)

3. From Alkynes (Hydration)

Terminal alkynes give aldehydes; internal alkynes give ketones.

\( RC{\equiv}CH + H_2O \xrightarrow{Hg^{2+}, H_2SO_4} RCHO \)
\( RC{\equiv}CR' + H_2O \rightarrow RCOR' \)

4. From Acid Chlorides

Rosenmund’s Reduction: Acid chlorides to aldehydes using hydrogen over Pd/BaSO₄.

\( RCOCl + H_2 \xrightarrow{Pd/BaSO_4} RCHO + HCl \)

Reaction with Dialkyl Cadmium: Acid chlorides react with dialkyl cadmium to form ketones.

\( RCOCl + (R')_2Cd \rightarrow RCOR' \)

5. From Calcium Salt of Carboxylic Acids

Dry distillation of calcium salts of fatty acids gives ketones.

\( (RCOO)_2Ca \xrightarrow{\Delta} RCOR + CaCO_3 \)

6. From Nitriles

Stephen Reaction: Nitriles are reduced to imines, then hydrolyzed to aldehydes.

\( RCN + SnCl_2/HCl \rightarrow RCH=NH \xrightarrow{H_2O} RCHO \)

Grignard’s Reagent: Nitriles react with RMgX and then are hydrolyzed to ketones.

\( RCN + R'MgX \xrightarrow{H_3O^+} RCOR' \)

7. From Esters

Grignard reagent reacts with esters to form ketones (intermediate stage before tertiary alcohol).

\( RCOOR' + R''MgX \xrightarrow{dry~ether} RCOR'' \)

Aldehydes and Ketones: Properties and Reactions

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Aldehydes and Ketones: Properties and Reactions

Physical Properties

State and Boiling Point: Aldehydes and ketones are typically liquids at room temperature. Their boiling points are higher than alkanes but lower than alcohols due to the presence of the polar carbonyl group.

Chemical Properties

Nucleophilic Addition Reactions

Ammonia and its Derivatives: React with aldehydes and ketones to form imines.
Hydrogen Cyanide (HCN): Adds to the carbonyl group to form cyanohydrins.
Sodium Bisulfite (NaHSO₃): Forms adducts with aldehydes and ketones.
Grignard Reagents: React to form alcohols.

Oxidation Reactions

Aldehydes: Oxidize to carboxylic acids.
Ketones: Generally resistant to oxidation.
Iodoform Test: Methyl ketones react with iodine and NaOH to form iodoform (CHI₃), a yellow precipitate.

Reduction Reactions

To Alcohols: Aldehydes reduce to primary alcohols; ketones to secondary alcohols using reducing agents like NaBH₄ or LiAlH₄.
To Alkanes: Clemmensen reduction (Zn/HCl) and Wolff–Löffler reduction (NH₂NH₂/KOH) reduce aldehydes and ketones to alkanes.

Base-Catalyzed Reactions

Aldol Condensation: Aldehydes and ketones with α-hydrogens undergo aldol condensation in the presence of a base, forming β-hydroxy aldehydes or ketones.
Cross Aldol Condensation: Involves two different aldehydes or ketones.
Cannizzaro Reaction: Aldehydes without α-hydrogens undergo disproportionation in the presence of a strong base, forming an alcohol and a carboxylic acid.

Tests for Aldehydes and Ketones

Tollens' Test: Aldehydes react with Tollens' reagent (AgNO₃/NH₃) to form a silver mirror. Ketones do not respond to this test.
Fehling's Test: Aldehydes react with Fehling's solution to form a red precipitate of Cu₂O. Ketones do not react.
Iodoform Test: Methyl ketones react with iodine and NaOH to form a yellow iodoform precipitate. Aldehydes do not give this test.

Uses of Aldehydes and Ketones

Aldehydes

Formaldehyde: Used in disinfectants, resins, and plastics.
Acetaldehyde: Used in the synthesis of acetic acid and perfumes.
Benzaldehyde: Used in perfumery and in dye industries.

Ketones

Acetone: Common solvent in nail polish removers and industrial cleaning.
Methyl Ethyl Ketone (MEK): Used in the production of synthetic rubber and plastics.
Cyclohexanone: Used in the manufacture of nylon.

Benzaldehyde: Properties and Reactions

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Benzaldehyde (C₆H₅CHO)

1. Lab Preparation from Toluene

Oxidation with Chromyl Chloride (Etard Reaction):

\( C_6H_5CH_3 + CrO_2Cl_2 \rightarrow C_6H_5CH(Cl)_2 \rightarrow C_6H_5CHO \)

This reaction is performed in a dry solvent like carbon tetrachloride, followed by hydrolysis to yield benzaldehyde.

2. Physical Properties

State: Colorless liquid with an almond-like odor.
Boiling Point: ~179°C
Moderately soluble in water; miscible with ethanol and ether.
Stable under normal conditions, but oxidizes in air to benzoic acid.

3. Chemical Properties

i. Oxidation

\( C_6H_5CHO + [O] \rightarrow C_6H_5COOH \) (Benzoic acid)

ii. Reduction

To benzyl alcohol: \( C_6H_5CHO + 2[H] \rightarrow C_6H_5CH_2OH \)
To toluene (Clemmensen or Wolff–Kishner): \( C_6H_5CHO \rightarrow C_6H_5CH_3 \)

iii. Nucleophilic Addition

With HCN: \( C_6H_5CHO + HCN \rightarrow C_6H_5CH(OH)CN \)
With NaHSO₃: \( C_6H_5CHO + NaHSO_3 \rightarrow C_6H_5CH(OH)SO_3Na \)

iv. With Ammonia and Derivatives

Hydroxylamine: → Oxime
Hydrazine: → Hydrazone
Phenylhydrazine: → Phenylhydrazone

v. With Phosphorus Pentachloride (PCl₅)

\( C_6H_5CHO + PCl_5 \rightarrow C_6H_5CHCl_2 + POCl_3 + HCl \)

vi. Cannizzaro Reaction

\( 2C_6H_5CHO + OH^- \rightarrow C_6H_5CH_2OH + C_6H_5COO^- \)

vii. Benzoin Condensation

In presence of CN⁻, two molecules of benzaldehyde condense to form benzoin.

\( 2C_6H_5CHO \xrightarrow{CN^-} C_6H_5CH(OH)COC_6H_5 \)

viii. Perkin’s Reaction

Reaction with acetic anhydride and sodium acetate to form cinnamic acid.

\( C_6H_5CHO + (CH_3CO)_2O \xrightarrow{NaOAc} C_6H_5CH=CHCOOH \)

ix. Electrophilic Substitution

Halogenation (o-/p-): with Br₂ in acetic acid
Nitration: \( C_6H_5CHO + HNO_3 \rightarrow o/p \)-nitrobenzaldehyde
Sulphonation: with concentrated H₂SO₄

4. Test: Aromatic vs Aliphatic Aldehydes

Fehling’s Test: Benzaldehyde does not give a positive test (no red ppt), while aliphatic aldehydes do.
Tollens' Test: Both aromatic and aliphatic aldehydes give silver mirror.

5. Uses of Benzaldehyde

Flavoring agent in food industry (almond essence)
Intermediate in dye, pharmaceutical, and perfume production
Used in synthesis of cinnamic acid, benzoin, and phenylacetic acid

Carboxylic Acids - Notes

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Carboxylic Acids

1. Classification

Monocarboxylic acids: One –COOH group
Example: Formic acid (HCOOH), Acetic acid (CH₃COOH)
Dicarboxylic acids: Two –COOH groups
Example: Oxalic acid (HOOC–COOH), Malonic acid (CH₂(COOH)₂)

2. General Formula and Structure

General formula: \( R{-}COOH \)

The carboxyl group consists of a carbon double bonded to an oxygen and single bonded to a hydroxyl group.

\( \ce{R-C(=O)OH} \)

3. Nomenclature

Common names: Derived from source (e.g. formic acid from ants).
IUPAC: Replace -e of alkane with -oic acid.
Example: Methanoic acid (HCOOH), Ethanoic acid (CH₃COOH)

4. Acidic Nature

Carboxylic acids are weak acids due to resonance stabilization of the carboxylate ion:

\( \ce{R-COOH R-COO^- + H^+} \)

Acidity is affected by the inductive effect and hydrogen bonding.

5. Methods of Preparation

From Alcohols:
\( \ce{RCH_2OH + [O] -> RCOOH} \)
From Aldehydes:
\( \ce{RCHO + [O] -> RCOOH} \)
From Nitriles:
\( \ce{RCN + 2H_2O + H^+ -> RCOOH + NH_3} \)
From Grignard's Reagent:
\( \ce{RMgX + CO_2 -> RCOOMgX ->[H^+] RCOOH} \)

6. Physical Properties

Carboxylic acids are generally liquids or solids at room temperature.
Lower members are miscible with water due to hydrogen bonding.
They have high boiling points compared to corresponding alcohols and ketones.

7. Chemical Properties

Acid-base reactions: Form salts with bases (e.g. NaOH)
Formation of esters: Reaction with alcohols in the presence of acid catalyst.
Reduction: To primary alcohols using LiAlH₄
Decarboxylation: Loss of CO₂ under heating with soda lime

8. Uses of Carboxylic Acids

Formic acid: Used in dyeing and textile finishing.
Acetic acid: Used as a food preservative and industrial solvent.
Benzoic acid: Used as a food preservative (E210).
Oxalic acid: Used in cleaning and bleaching agents.

Carboxylic Acids: Properties and Reactions

1. Physical Properties

State: Colorless liquids or solids.
Boiling Point: Varies; e.g., acetic acid boils at 118°C.
Solubility: Soluble in water; solubility decreases with increasing molecular size.

2. Chemical Properties

i. Acidic Character

Carboxylic acids are weak acids. The presence of electron-withdrawing groups increases their acidity, while electron-donating groups decrease it.

CH₃COOH ⇌ CH₃COO⁻ + H⁺

ii. Reactions with Active Metals

2CH₃COOH + 2Na → 2CH₃COONa + H₂↑

iii. Reactions with Alkalies

CH₃COOH + NaOH → CH₃COONa + H₂O

iv. Reactions with Carbonates and Bicarbonates

2CH₃COOH + Na₂CO₃ → 2CH₃COONa + H₂O + CO₂↑

v. Formation of Acid Derivatives

Carboxylic acids react with alcohols to form esters in a reaction known as esterification.

CH₃COOH + CH₃OH → CH₃COOCH₃ + H₂O

vi. Decarboxylation

Carboxylic acids can lose a CO₂ molecule when heated with soda lime (NaOH + CaO) to form alkanes.

CH₃COOH → CH₄ + CO₂

vii. Hell-Volhard-Zelinsky (HVZ) Reaction

Carboxylic acids undergo α-halogenation in the presence of halogens (Cl₂ or Br₂) and red phosphorus.

RCOOH + X₂ → α-Halo acid

viii. Electrophilic Substitution in Aromatic Carboxylic Acids

The carboxyl group is meta-directing and deactivates the benzene ring towards electrophilic substitution.

Nitration: C₆H₅COOH + HNO₃ → m-nitrobenzoic acid
Halogenation: C₆H₅COOH + Cl₂ → m-chlorobenzoic acid
Sulphonation: C₆H₅COOH + H₂SO₄ → m-benzoic acid

3. Tests for Carboxylic Acids

Reaction with Sodium Bicarbonate: Formation of CO₂ gas indicates the presence of a carboxyl group.
Reaction with Alcohols: Formation of esters upon heating with alcohols and acid catalyst.
Reaction with Thionyl Chloride (SOCl₂): Formation of acyl chlorides.

4. Uses of Carboxylic Acids

Formic Acid (HCOOH): Used in leather industry, as a preservative, and in textile finishing.
Acetic Acid (CH₃COOH): Used in food industry as vinegar, in manufacturing of synthetic fibers, and as a solvent.
Benzoic Acid (C₆H₅COOH): Used as a food preservative, in manufacturing of dyes, and as an intermediate in chemical synthesis.

9. Organic compounds containing Nitrogen

Aliphatic Amines - Notes

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Aliphatic Amines

1. General Formula and Classification

General formula: \( R{-}NH_2 \) (for primary amines)

Primary (1°): One alkyl group bonded to nitrogen
\( R{-}NH_2 \)
Secondary (2°): Two alkyl groups bonded to nitrogen
\( R_2{-}NH \)
Tertiary (3°): Three alkyl groups bonded to nitrogen
\( R_3{-}N \)

2. Structure of the Amino Group

The nitrogen atom in amines is \( sp^3 \)-hybridized with a lone pair, giving a trigonal pyramidal geometry.

3. Nomenclature

Common system: Alkylamine (e.g., methylamine)
IUPAC: Alkanamine (e.g., methanamine, ethanamine)

4. Methods of Preparation

Reduction of nitro compounds:
\( \ce{R{-}NO_2 ->[H_2,Ni] R{-}NH_2} \)
Ammonolysis of alkyl halides:
\( \ce{R{-}X + NH_3 -> R{-}NH_2 + HX} \)
Reduction of nitriles:
\( \ce{RCN ->[H_2,Ni] RCH_2NH_2} \)
Gabriel phthalimide synthesis:
Useful for preparation of 1° amines.
Hofmann bromamide reaction:
\( \ce{RCONH_2 ->[Br_2,OH^-] RNH_2} \)

5. Physical Properties

Lower amines are gases; higher ones are liquids or solids.
They form hydrogen bonds — 1° > 2° > 3° in boiling points.
Soluble in water (especially lower amines)

6. Chemical Properties

Basic nature: Amines accept protons due to lone pair on nitrogen.
\( \ce{R{-}NH_2 + HCl -> R{-}NH_3^+Cl^-} \)
Acylation: With acid chlorides or anhydrides (1° and 2° amines).
Alkylation: Further substitution with alkyl halides.
Reaction with nitrous acid:
- 1° amine: forms alcohol with N₂ evolution
- 2° amine: forms nitrosoamine (yellow oily liquid)
- 3° amine: no reaction or forms ammonium salt
Carbylamine reaction: Only 1° amines form isocyanides.
\( \ce{RNH_2 + CHCl_3 + KOH -> RNC + 3KCl + 3H_2O} \)

7. Identification of 1°, 2°, 3° Amines

Hinsberg’s Test: Distinguishes between 1°, 2°, and 3° amines.
Carbylamine Test: Positive only for 1° amines.
Reaction with nitrous acid: Evolving gas (N₂) indicates 1° amine.

8. Uses of Amines

Used in manufacture of dyes, drugs, and polymers.
Methylamine and ethylamine are used as solvents and intermediates.
Aniline is a key intermediate in rubber processing and dye production.

Amines - Classification and Preparation

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Amines

1. Nomenclature

Common names: Alkyl + "amine" (e.g. methylamine)
IUPAC: Replace final ‘e’ in alkane with ‘amine’ (e.g. methanamine)

2. Classification of Amines

Primary (1°): One alkyl/aryl group attached to nitrogen
\( \ce{RNH2} \)
Secondary (2°): Two alkyl/aryl groups
\( \ce{R2NH} \)
Tertiary (3°): Three alkyl/aryl groups
\( \ce{R3N} \)

3. Structure & General Formula

Amines contain a nitrogen atom with a lone pair, making them nucleophilic.

General formula (for aliphatic primary amines): \( \ce{C_nH_{2n+1}NH_2} \)

4. Methods of Preparation

i) From Alcohol

Reduction of alcohol to amine (via alkyl halide or nitrile):

\( \ce{R{-}OH ->[PCl_5] R{-}Cl ->[NH_3] R{-}NH_2} \)

ii) From Alkyl Halide

Alkyl halide reacts with excess ammonia (ammonolysis):

\( \ce{R{-}X + 2NH_3 -> RNH_2 + NH_4X} \)

iii) From Nitrile

Reduction of nitrile using hydrogen or lithium aluminium hydride:

\( \ce{RCN ->[H_2, Ni] RCH_2NH_2} \)

iv) From Amide (Hofmann’s Degradation)

Amide reacts with bromine and alkali to form a primary amine with one carbon less:

\( \ce{RCONH_2 + Br_2 + 4NaOH -> RNH_2 + Na_2CO_3 + 2NaBr + 2H_2O} \)

v) From Nitro Compounds

Reduction of nitroalkane or nitrobenzene to amine:

\( \ce{RNO_2 ->[Sn/HCl] RNH_2} \)

vi) Gabriel Phthalimide Synthesis

Used to prepare primary amines (especially aliphatic):

\( \ce{Phthalimide + R{-}X ->[KOH] RNH_2 + Phthalic~acid} \)

Amines - Properties

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Amines: Physical and Chemical Properties

1. Physical Properties

State:
- Lower amines (C1–C3): Gaseous
- Higher ones: Liquids or solids
Solubility: 1° and 2° amines form hydrogen bonds with water, hence more soluble than 3°.
Boiling Point:
- 1° > 2° > 3° (due to decreasing hydrogen bonding)
- Alcohols > Amines (due to more electronegative O in OH)

2. Basic Character of Amines

Basicity is due to lone pair of electrons on nitrogen.

Order in aqueous solution:

2° > 1° > 3° > NH₃ (in gas phase, 3° is more basic)
Aniline is less basic than alkyl amines due to resonance delocalization of lone pair on N.

Effect of Substituents on Aniline's Basicity

Electron withdrawing groups (e.g., NO₂) decrease basicity.
Electron donating groups (e.g., CH₃, OH) increase basicity.

3. Alkylation (Mechanism)

Amines react with alkyl halides to form higher amines and quaternary salts:

\( \ce{RNH_2 + R'X -> RR'NH + HX} \)

Excess alkyl halide can lead to quaternary ammonium salt:

\( \ce{R_3N + R'X -> R_4N^+X^-} \)

4. Acylation (Mechanism)

1° and 2° amines react with acid chlorides or anhydrides:

\( \ce{RNH_2 + R'COCl -> RNHCO{-}R' + HCl} \)

Mechanism involves nucleophilic attack of lone pair on carbonyl carbon.

5. Reaction with Nitrous Acid

1° amine: forms alcohol and evolves N₂ gas
2° amine: forms N-nitrosoamine
3° amine: no visible reaction

\( \ce{RNH_2 + HNO_2 -> ROH + N_2 + H_2O} \)

6. Carbylamine Reaction

Used to test for 1° amines (foul-smelling isocyanide formed):

\( \ce{RNH_2 + CHCl_3 + 3KOH -> RNC + 3KCl + 3H_2O} \)

7. Distinction Between 1°, 2°, and 3° Amines (Hinsberg’s Test)

1° amine: Forms soluble sulfonamide (gives precipitate with HCl)
2° amine: Forms insoluble sulfonamide
3° amine: No reaction with Hinsberg’s reagent

Aniline - Preparation and Properties

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Aniline: Preparation and Properties

1. Preparation

By reduction of nitrobenzene:

\( \ce{C6H5NO2 + 3H2 -> C6H5NH2 + 2H2O} \)
(Using Sn/HCl or Fe/HCl as reducing agents)

2. Physical Properties

State: Liquid (oil-like)
Solubility: Slightly soluble in water; soluble in organic solvents
Boiling Point: 184 °C (higher due to hydrogen bonding)

3. Chemical Properties

a) Reaction with acids

Aniline forms salts with mineral acids like HCl and H₂SO₄:

\( \ce{C6H5NH2 + HCl -> C6H5NH3^+Cl^-} \)

b) Acetylation

Reacts with acetic anhydride to form acetanilide (less reactive amide):

\( \ce{C6H5NH2 + (CH3CO)2O -> C6H5NHCOCH3 + CH3COOH} \)

c) Alkylation

Reaction with alkyl halides to form alkylated anilines:

\( \ce{C6H5NH2 + R-X -> C6H5NHR + HX} \)

d) Benzoylation

Reaction with benzoyl chloride:

\( \ce{C6H5NH2 + C6H5COCl -> C6H5NHCOC6H5 + HCl} \)

e) Carbylamine Reaction

Formation of isocyanide (test for primary amines):

\( \ce{C6H5NH2 + CHCl3 + 3KOH -> C6H5NC + 3KCl + 3H2O} \)

f) Diazotisation

Reaction with nitrous acid to form diazonium salt:

\( \ce{C6H5NH2 + HNO2 + HCl -> C6H5N2^+Cl^- + 2H2O} \)

g) Electrophilic substitution reactions

Bromination: Aniline reacts readily with bromine giving 2,4,6-tribromoaniline.
Nitration: Forms mainly 2-nitroaniline.
Sulphonation: Forms sulphonated derivatives (e.g., sulphanilic acid).

4. Tests for Aniline

Carbylamine test: Formation of foul-smelling isocyanide.
Diazotisation test: Formation of diazonium salt (used for coupling reactions).

5. Uses of Aniline

Manufacture of dyes (azo dyes).
Production of pharmaceuticals, rubber processing chemicals.
Intermediate in synthesis of pesticides, plastics.

6. Structural Formula of Aniline

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Cyanides and Isocyanides - Preparation Methods

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Cyanides and Isocyanides

Cyanides (R–CN)

Methods of Preparation

From Alkyl Halide:
\( R{-}X + KCN \rightarrow R{-}CN + KX \)
(via nucleophilic substitution using potassium cyanide)
From Amide:
\( RCONH_2 \xrightarrow{P_2O_5} R{-}CN + H_2O \)
(dehydration of amide using phosphorus pentoxide)

Isocyanides (R–NC)

Methods of Preparation

From Alkyl Halide:
\( R{-}X + AgNC \rightarrow R{-}NC + AgX \)
(reaction with silver isocyanide)
From Primary Amines:
\( RNH_2 + CHCl_3 + 3KOH \rightarrow R{-}NC + 3KCl + 3H_2O \)
(Carbylamine reaction)

Diazonium Salts - Preparation and Reactions

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Diazonium Salts

Preparation

Diazonium salts are prepared by treating **primary aromatic amines** with **nitrous acid** at low temperatures (0–5 °C).

From Aniline:

\[
C_6H_5NH_2 + HNO_2 + HCl \rightarrow C_6H_5N_2^+Cl^- + 2H_2O
\]

(Nitrous acid is generated in situ from sodium nitrite and hydrochloric acid.)

Properties and Reactions

1. Sandmeyer’s Reaction

Used to replace the diazonium group with halogens.

\[
C_6H_5N_2^+Cl^- + CuCl \rightarrow C_6H_5Cl + N_2 \uparrow
\]

2. Gattermann Reaction

Similar to Sandmeyer but uses Cu powder and HCl.

\[
C_6H_5N_2^+Cl^- + Cu + HCl \rightarrow C_6H_5Cl + N_2 + CuCl
\]

3. Balz–Schiemann Reaction

Used to replace the diazonium group with fluorine.

\[
C_6H_5N_2^+BF_4^- \xrightarrow{\Delta} C_6H_5F + BF_3 + N_2
\]

4. Replacement Reactions

By Hydrogen: \( C_6H_5N_2^+Cl^- + H_3PO_2 \rightarrow C_6H_6 + N_2 \uparrow \)
By Hydroxyl: \( C_6H_5N_2^+ + H_2O \xrightarrow{\Delta} C_6H_5OH + N_2 \)
By Nitro: \( C_6H_5N_2^+ + CuNO_2 \rightarrow C_6H_5NO_2 + N_2 \)

5. Coupling Reactions

Diazonium salts undergo electrophilic substitution with phenols and aromatic amines to form azo compounds.

With Phenol:
\[
C_6H_5N_2^+ + C_6H_5OH \rightarrow C_6H_5{-}N{=}N{-}C_6H_4OH
\]
With Aniline:
\[
C_6H_5N_2^+ + C_6H_5NH_2 \rightarrow C_6H_5{-}N{=}N{-}C_6H_4NH_2
\]

Importance in Organic Synthesis

Key intermediates in the preparation of aromatic halides, phenols, and azo dyes.
Useful in synthetic routes that avoid direct substitution on the benzene ring.

10. Biomolecules

Carbohydrates - Biology Notes

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Carbohydrates – Study Notes

Definition

Carbohydrates are organic compounds composed of carbon, hydrogen, and oxygen in the ratio C_n(H₂O)_n. They are also called saccharides.

Classification

D- and L- Configuration

D and L refer to the configuration around the asymmetric carbon farthest from the carbonyl group, based on glyceraldehyde.

Structural Formulas (Glucose and Fructose)

Reactions

Test for Glucose and Fructose

Bromine Water Test (for Glucose):

\[ \ce{C6H12O6 + Br2 + H2O -> C6H12O7 + 2HBr} \]

Fructose does not react directly with bromine water.

Disaccharides

Sucrose: Glucose + Fructose (non-reducing)
Maltose: Glucose + Glucose (reducing)
Lactose: Glucose + Galactose (reducing)
Linkage: Glycosidic bond (e.g., α(1→4) in maltose)

Polysaccharides

Starch: Energy store in plants, α(1→4) and α(1→6) bonds.
Cellulose: Structural, β(1→4) glucose linkages (not digestible).
Glycogen: Energy storage in animals, highly branched.

Importance of Carbohydrates

Primary energy source (e.g., glucose).
Structural components (e.g., cellulose in plants).
Biochemical roles (cell recognition, immune response).

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Proteins - Biology Notes

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Proteins – Study Notes

Structural Units of Proteins

Proteins are polymers made up of amino acids linked by peptide bonds forming polypeptides.

Amino Acids – General Structure

General formula of an amino acid:

\[
\ce{NH2-CH(R)-COOH}
\]

where R is the variable side chain.

Amino acids exist as zwitterions at physiological pH:

\[
\ce{^{+}H3N-CH(R)-COO^{-}}
\]

The pH at which the amino acid has no net charge is called the isoelectric point (pI).

Classification of Proteins

Based on molecular shape: fibrous, globular, and membrane proteins.
Protein Structures:
- Primary: sequence of amino acids.
- Secondary: α-helix, β-pleated sheets (hydrogen bonding).
- Tertiary: overall 3D folding due to interactions between side chains.
- Quaternary: assembly of multiple polypeptide chains.
Denaturation: loss of native structure due to heat, pH, chemicals.

Enzymes and Hormones

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Enzymes

Biological catalysts that speed up chemical reactions without being consumed.

Mostly proteins with specific active sites for substrate binding.

Hormones

Chemical messengers that regulate physiology and behavior.

Examples include insulin (protein hormone) and steroid hormones.

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Vitamins - Biology Notes

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Vitamins – Study Notes

Classification of Vitamins

Fat-soluble vitamins: A, D, E, K
Water-soluble vitamins: B-complex group and C

List of Vitamins

Vitamin A: Important for vision and growth.
Vitamin B complex: Group of vitamins involved in metabolism (energy release, nerve function).
Vitamin C: Antioxidant, helps in collagen formation.
Vitamin D: Regulates calcium and phosphate balance, bone health.
Vitamin E: Antioxidant protecting cells from damage.
Vitamin K: Important for blood clotting.

Functions and Deficiency Diseases

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Vitamin Functions and Deficiency Diseases

Vitamin A: Maintains healthy vision, skin, and immune system.
Deficiency: Night blindness, xerophthalmia.
Vitamin B complex: Supports energy metabolism and nervous system.
Deficiency: Beriberi (B1), pellagra (B3), anemia (B12).
Vitamin C: Antioxidant and aids wound healing.
Deficiency: Scurvy (bleeding gums, weakness).
Vitamin D: Promotes calcium absorption.
Deficiency: Rickets in children, osteomalacia in adults.
Vitamin E: Protects cells from oxidative damage.
Deficiency: Rare, may cause nerve and muscle damage.
Vitamin K: Essential for blood clotting.
Deficiency: Bleeding disorders.

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Nucleic Acids - Biology Notes

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Nucleic Acids – Study Notes

Basic Units

Nucleic acids are polymers made of nucleotides. Each nucleotide consists of:

Purines: Adenine (A), Guanine (G)
Pyrimidines: Cytosine (C), Thymine (T in DNA), Uracil (U in RNA)

DNA Structure

RNA Structure

RNA is usually single-stranded and contains ribose sugar instead of deoxyribose.

Contains uracil (U) instead of thymine.

No complex chemical structure required here as per syllabus.

Differences Between DNA and RNA

Feature	DNA	RNA
Strand	Double-stranded (double helix)	Single-stranded
Sugar	Deoxyribose	Ribose
Bases	A, T, G, C	A, U, G, C
Function	Stores genetic information	Protein synthesis, gene regulation
Location	Mostly nucleus	Nucleus and cytoplasm

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