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 %
Examples: Sugar or salt in water
Solubility: Increases with temperature for most solids.
Concentration: Given by:
\[
\text{Molarity (M)} = \frac{\text{moles of solute}}{\text{volume of solution in L}}
\]
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.
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.
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.
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:
2nd Law (Relative Lowering): The relative lowering of vapour pressure is equal to the mole fraction of the solute.
\[
\frac{P^0 - P}{P^0} = \chi_{\text{solute}} = \frac{n_B}{n_A + n_B}
\]
Boiling point of a solution is higher than that of pure solvent.
\[
\Delta T_b = K_b \cdot m
\]
Where:
Freezing point of a solution is lower than that of the pure solvent.
\[
\Delta T_f = K_f \cdot m
\]
Where:
Osmotic pressure is the pressure required to stop the flow of solvent through a semipermeable membrane into the solution.
\[
\Pi = C R T
\]
Where:
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:
Sometimes, calculated molar mass using colligative properties is different from the actual value. This is called abnormal molecular mass.
Occurs due to:
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:
For association:
\[
i = 1 - \alpha \left(1 - \frac{1}{n}\right)
\]
Where \( n \) is the number of molecules associating to form 1 unit.
To calculate molar mass in cases of dissociation or association:
\[
M = \frac{1000 \cdot K \cdot w_2}{\Delta T \cdot w_1 \cdot i}
\]
Definition: Molarity is the number of moles of solute per litre of solution.
\[
M = \frac{n}{V}
\]
Where:
Example: 1 mole of NaCl in 1 litre of solution → 1 M solution.
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.
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.
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
\[
\chi_{\text{ethanol}} = \frac{2}{2+3} = 0.4
\]
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:
\[
\text{ppm} = \frac{2}{1000000} \times 10^6 = 2 \, \text{ppm}
\]
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: 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:
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:
\[
C = \frac{P}{k_H} = \frac{0.21}{1.3 \times 10^5} \approx 1.615 \times 10^{-6} \, \text{mol/L}
\]
Q: The solubility of CO₂ in water at 1 atm is 0.030 mol/L. What will be the solubility at 2 atm?
Solution:
\[
C_1 = 0.030 \, \text{mol/L at } P_1 = 1 \, \text{atm}
\]
\[
C_2 = \frac{P_2}{P_1} \cdot C_1 = \frac{2}{1} \cdot 0.030 = 0.060 \, \text{mol/L}
\]
Raoult’s Law states that the partial vapor pressure of the solvent is directly proportional to its mole fraction in the solution.
\[
P_1 = X_1 P_1^0
\]
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
When both solute and solvent are volatile:
\[
P_{\text{total}} = P_1 + P_2 = X_1 P_1^0 + X_2 P_2^0
\]
Definition: A mixture of two liquids that boils at a constant temperature and behaves like a pure substance.
Types:
Graph of vapor pressure vs mole fraction:
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:
Colligative properties are used to find the molar mass \( M_2 \) of an unknown solute based on changes observed in a known solvent.
(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:
\[
\Pi = CRT = \frac{n_2}{V}RT = \frac{W_2 RT}{M_2 V}
\]
The more the depression/elevation/pressure change, the greater the number of solute particles—this allows for indirect determination of unknown molecular masses.
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.
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:
\[
\frac{P^0 - P_s}{P^0} = \frac{n_2}{n_1 + n_2} \approx \frac{n_2}{n_1}
\]
When \( n_2 \ll n_1 \), \( n_1 + n_2 \approx n_1 \), so:
\[
\frac{\Delta P}{P^0} = \frac{n_2}{n_1}
\]
We use Raoult’s Law to determine the relative molecular mass of a non-volatile solute:
\[
\frac{\Delta P}{P^0} = \frac{W_2 \cdot M_1}{W_1 \cdot M_2}
\]
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:
The depression in freezing point is directly proportional to the molality of the solution.
\[
\Delta T_f = K_f \cdot m
\]
where:
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:
\[
M_2 = \frac{K_f \cdot W_2}{\Delta T_f \cdot W_1}
\]
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.
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:
The elevation in boiling point is directly proportional to the molality of the solution:
\[
\Delta T_b = K_b \cdot m
\]
where:
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:
\[
M_2 = \frac{K_b \cdot W_2}{\Delta T_b \cdot W_1}
\]
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 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.
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.
Used to determine the relative molecular mass of solutes, especially macromolecules. The formula is:
π = CRT
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.
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
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.
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
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 = \(\dfrac{\text{Observed colligative property}}{\text{Calculated colligative property assuming no dissociation/association}}\)
The colligative property formulas are modified by multiplying with i:
If a solute dissociates into n particles, then:
i = 1 + (n − 1)α
If n molecules associate to form 1 unit:
i = 1 − α(n − 1)/n
Solution:
π = iCRT → 4.92 = i × 0.1 × 0.0821 × 298
i = 4.92 / (0.1 × 0.0821 × 298) ≈ 2.0
➤ Almost complete dissociation.
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.
π = C × R × T
π = 0.1 × 0.0821 × 298 ≈ 2.45 atm
π = iCRT = 2 × 0.2 × 0.0821 × 300 ≈ 9.85 atm
i = Normal / Observed = 60 / 30 = 2
Since i > 1, the solute undergoes dissociation.
i = 60 / 120 = 0.5
Since i < 1, the solute undergoes association.
i = 1 + (n − 1)α → 2.5 = 1 + 2α → α = 0.75
Degree of dissociation is 75%.
i = 1 − α(n − 1)/n → 0.6 = 1 − α(1)/2 → α = 0.8
Degree of association is 80%.
ΔTb = i × Kb × m = 2 × 0.52 × 0.1 = 0.104 °C
ΔTf = i × Kf × m = 1 × 1.86 × 0.2 = 0.372 °C
Freezing point = 0 − 0.372 = −0.372 °C
Moles of NaOH = 40 / 40 = 1 mol
Moles of H2O = 180 / 18 = 10 mol
χNaOH = 1 / (1 + 10) = 1/11 ≈ 0.0909
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
ΔTf = i × Kf × m → 0.372 = i × 1.86 × 0.1
i = 0.372 / (1.86 × 0.1) = 2.0 (full dissociation)
i = Normal / Observed = 60 / 120 = 0.5 (indicates dimerization)
π = iCRT = 1 × 0.2 × 0.0821 × 298 = 4.89 atm
ΔTb = i × Kb × m = 2 × 0.52 × 1 = 1.04 °C
ΔTf = i × Kf × m = 1 × 1.86 × 0.1 = 0.186 °C
Tf = 0 − 0.186 = −0.186 °C
Oxidation: Loss of electrons (at the anode)
Reduction: Gain of electrons (at the cathode)
Electron flow: From anode → cathode
EMF is the potential difference between two electrodes when no current is flowing.
EMF = Ecathode − Eanode
Measured in volts (V)
The potential of a half-cell under standard conditions (1 M, 1 atm, 25°C).
Relates electrode potential to concentration/pressure of species involved.
E = E° − (0.0591 / n) log Q
Where:
Relates cell EMF to the spontaneity of the reaction.
ΔG = −nFE
If ΔG < 0 → reaction is spontaneous (E > 0)
Statement: At infinite dilution, molar conductivity is the sum of individual ion conductivities:
Λ∞m
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.
Redox stands for reduction-oxidation reactions, where electrons are transferred:
The flow of electrons occurs from the anode to the cathode through the external circuit.
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.
A Galvanic cell is written in the following format:
Double line (||) represents the salt bridge.
Galvanic cells are based on redox reactions:
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).
Definition: SHE is a reference electrode with a potential of 0.00 V under standard conditions.
Construction:
Definition: Electrode potential measured under standard conditions: 1 M solution, 1 atm gas, 25°C (298 K).
Measured using: SHE as reference.
Cell setup:
Standard cell notation follows:
The Electrochemical Series is a list of standard reduction potentials (E°) of various elements and ions arranged in decreasing order.
Measured under standard conditions: 1 M concentration, 1 atm gas pressure, and 25°C (298 K).
Higher E° → greater tendency to gain electrons (reduction)
A redox reaction is spontaneous if the overall cell potential (E°cell) is positive:
The Nernst Equation relates the electrode potential under non-standard conditions to standard electrode potential.
Where:
The free energy change (ΔG) of a redox reaction is related to the EMF of the cell:
Where:
At equilibrium, the free energy is related to the equilibrium constant (K) as:
Also:
E° = (0.0591 / n) × log K
To determine whether a redox reaction is spontaneous:
Conductance (G) is the reciprocal of resistance (R):
Unit of resistance: ohm (Ω)
Unit of conductance: siemens (S) or mho
Unit:
- Resistivity (ρ): Ω·cm
- Conductivity (κ): S·cm⁻¹
The cell constant (denoted by K) relates the physical dimensions of a conductivity cell:
To calculate: Use a standard solution (like KCl) of known conductivity (κ):
It is the conductance of all the ions produced by one gram-equivalent of an electrolyte in a given volume.
Units: S·cm²·equiv⁻¹
Molar conductance is the conductance of all ions produced by one mole of electrolyte in solution.
Units: S·cm²·mol⁻¹
Graphical Representation:
| 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 | Λm | S·cm²·mol⁻¹ |
| Equivalent Conductance | Λeq | S·cm²·equiv⁻¹ |
Given:
Given:
Given:
Definition: Conductance of all ions produced by one mole of an electrolyte in a given volume of solution.
At infinite dilution, electrolytes are completely dissociated. The conductance under this condition is denoted by:
Note: For weak electrolytes like CH₃COOH, Λm∞ is determined using **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.*
Using Kohlrausch’s Law:
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:
\[
m = Z \cdot Q = Z \cdot I \cdot t
\]
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:
\[
\frac{m_1}{m_2} = \frac{E_1}{E_2}
\]
The relationship is given by:
\[
F = N_A \cdot e
\]
Calculation:
\[
F = (6.022 \times 10^{23}) \times (1.602 \times 10^{-19}) \approx \boxed{96485 \, \text{C/mol}}
\]
Use: Flashlights, clocks, remote controls
Use: Hearing aids, watches, calculators
Charging: Reverse of discharging reactions
Use: Inverters, automobiles, backup power
Overall Reaction:
Use: Spacecrafts, submarines, clean energy systems
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:
In moist air, iron acts as an anode and undergoes oxidation, while oxygen acts as the cathode and is reduced.
The Fe²⁺ ions react with water and oxygen to form hydrated ferric oxide (rust).
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).
Average Rate:
Instantaneous Rate:
Graphical Representation: Slope of concentration vs. time graph gives the rate at a point.
Order: Sum of powers of concentration in the rate law.
Molecularity: Number of reacting species involved in an elementary step (1, 2, or 3).
Rate Law: Expresses rate as a function of reactant concentrations.
Specific Rate Constant (k): The rate when all reactants have unit concentration.
Reactions occur due to collisions between reacting particles. For a reaction to occur:
Activation Energy (Ea): Minimum energy needed for a reaction to occur.
Threshold Energy: Total energy required by reactants to convert into products.
\[
\text{Threshold Energy} = \text{Activation Energy} + \text{Energy of Reactants}
\]
Relates rate constant \( k \) with temperature and activation energy:
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:
Reactions are classified based on the time they take:
Explanation:
Reaction rate depends on the nature and strength of chemical bonds:
Example:
The rate of a chemical reaction is the change in concentration of a reactant or product per unit time.
The negative sign for reactants indicates a decrease in concentration over time.
For a general reaction:
The rate is given by:
By plotting concentration vs. time:
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."
Active mass refers to the effective concentration of a substance that takes part in a chemical reaction.
For a general reaction:
The rate of reaction is given by:
Consider the reaction:
According to the 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.
The rate law is an experimentally determined expression that relates the rate of a reaction to the concentrations of its reactants.
If more than one reactant is involved, the equation becomes:
The rate of reaction is directly proportional to the product of the concentrations of reactants, each raised to their respective orders:
At constant temperature:
This relation can be used to determine the order \( n \) experimentally by comparing rates at different concentrations.
The order of a chemical reaction is the sum of the powers of concentration terms in the experimentally determined rate law.
The order of a reaction is not always equal to the stoichiometric coefficients. It is determined experimentally.
Example:
Order is determined using initial rate data. Compare rates at different concentrations to find the order \( n \):
Rate is independent of concentration.
Starting with:
The time required for the concentration of a reactant to become half of its initial value.
Note: For a first-order reaction, half-life is independent of initial concentration.
Q: A first order reaction has a rate constant \( k = 1.386 \times 10^{-3} \, \text{s}^{-1} \). What is the half-life?
Solution:
Q: The concentration of a reactant drops from 1.0 M to 0.25 M in 2 hours. Find rate constant.
Solution:
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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.
For a simple (elementary) reaction:
Here:
So, for elementary reactions:
But for complex (multi-step) reactions, order may be different and must be determined experimentally.
| 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 |
Unimolecular reaction:
Bimolecular reaction:
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 (\(E_a\)).
During the reaction, reactant molecules form an unstable intermediate with high energy called the activated complex (or transition state).
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.
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
For a collision to result in a chemical reaction, it must be effective.
Conditions for effectiveness:
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.
A reaction mechanism is the step-by-step sequence of elementary reactions by which an overall chemical change occurs.
In a multi-step reaction, the slowest elementary step is called the rate-determining step.
This step controls the overall rate of the reaction.
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:
Rate law from slow step:
Order: 2 (from rate-determining step)
The unit of rate constant \( k \) depends on the order of reaction.
The rate constant \( k \) increases with temperature, described by the Arrhenius equation:
Where:
Taking natural logarithm on both sides:
In base-10 logarithm:
Plotting \( \log k \) vs \( \frac{1}{T} \) gives a straight line with:
If rate constants \( k_1 \) and \( k_2 \) are known at two temperatures \( T_1 \) and \( T_2 \), we use:
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:
Solve for \( E_a \) after calculating both sides.
- d-Block: Groups 3 to 12 (transition elements)
- f-Block: Lanthanides and Actinides, placed separately at the bottom
General Configuration of d-block:
- 3d series: Sc (Z = 21) to Zn (Z = 30)
- Elements are generally found in ores as oxides or sulfides
Preparation: From chromite ore by fusion and acidification
Important Reaction:
Acidification: \( 2Na_2CrO_4 + H_2SO_4 \rightarrow \text{K}_2\text{Cr}_2\text{O}_7 + H_2O \)
Properties:
Preparation: By oxidation of \( \text{MnO}_2 \) with KOH and O2 or KNO3
Properties:
Electronic configuration: General form:
Key trends:
Atomic Numbers: 57 (La) to 71 (Lu)
General configuration:
Properties:
Atomic Numbers: 89 (Ac) to 103 (Lr)
General configuration:
Properties:
| 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 |
General configuration:
Definition: Steady decrease in ionic radii from Ce to Lu
Cause: Poor shielding of 4f electrons
Consequences:
| 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) |
Step 1: Conversion of MnO₂ to Potassium Manganate (K₂MnO₄)
Step 2: Oxidation of manganate to permanganate
In Acidic Medium: Mn⁷⁺ reduces to Mn²⁺
In Basic Medium: MnO₄⁻ reduces to MnO₂
In Neutral Medium: MnO₄⁻ still reduces to MnO₂
Example: Redox titration of FeSO₄ with KMnO₄ in acid medium
Step 1: Fusion of chromite with Na₂CO₃ and O₂ to form sodium chromate
Step 2: Conversion of sodium chromate to sodium dichromate
\[
Na_2Cr_2O_7 + 2KCl \rightarrow K_2Cr_2O_7 + 2NaCl
\]
In Acidic Medium: Cr₂O₇²⁻ reduces to Cr³⁺
In Basic Medium: Converted to chromate ion (CrO₄²⁻) and acts less effectively
In Neutral Medium: Less reactive, requires catalyst or heat for oxidation
Example:
In Acidic Medium:
In Basic Medium:
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.
Ligands are atoms, ions, or molecules that donate a lone pair of electrons to the central metal ion.
Number of ligand atoms directly bonded to the central atom/ion.
Example: In \([Co(NH_3)_6]^{3+}\), the coordination number is 6.
It is calculated by treating all ligands as neutral or charged, summing their contributions, and solving for the metal's charge.
Example: In \([Fe(CN)_6]^{4-}\)
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.
Occurs due to differences in the structure or connectivity of ligands.
Example: \(\text{NO}_2^-\) can bind via N (nitro) or O (nitrito).
Occurs due to different spatial arrangements of ligands.
| 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 explains the bonding in coordination compounds based on the hybridization of orbitals in the central metal ion. It helps predict geometry and magnetic properties.
Uses inner (n−1)d orbitals for hybridization (d²sp³).
Example: \([Fe(CN)_6]^{4-}\)
Uses outer nd orbitals for hybridization (sp³d²).
Example: \([FeF_6]^{3-}\)
| 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 (CFT) explains the bonding and properties of coordination compounds by treating metal-ligand interactions as purely electrostatic.
In an octahedral field, the five degenerate d-orbitals of the metal split into two sets:
Examples:
In a tetrahedral field, the d-orbitals split in the opposite pattern:
Since \(\Delta_t\) is smaller than \(\Delta_0\), tetrahedral complexes are usually high spin (no pairing).
Color arises from the d–d electronic transitions between t2g and eg levels.
Example: \([Ti(H_2O)_6]^{3+}\) appears violet due to absorption in the green-yellow region.
| 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 |
The stability of a coordination compound refers to its ability to remain intact without decomposition. There are two types:
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 \([ML_n]\):
A higher \( K_f \) value indicates greater stability of the complex.
When ligands are added one at a time, we define stepwise stability constants \( K_1, K_2, \dots \)
\( \beta_n \) is the overall stability constant of the complex.
Coordination compounds are widely used in qualitative and quantitative analysis:
Many biological molecules are coordination compounds:
Coordination compounds play vital roles in:
Haloalkanes have the general formula: \( C_nH_{2n+1}X \), where \( X \) is a halogen (F, Cl, Br, I).
They undergo nucleophilic substitution:
\[
R{-}X + OH^- \rightarrow R{-}OH + X^-
\]
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.
Haloarenes undergo electrophilic substitution rather than nucleophilic substitution.
The halogen atom is deactivating but ortho-para directing due to resonance.
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.
Addition of hydrogen halide (HX) to an alkene:
\[
CH_2=CH_2 + HBr \rightarrow CH_3CH_2Br
\]
Follows Markovnikov's rule:
\[
CH_3CH=CH_2 + HCl \rightarrow CH_3CHClCH_3
\]
Alcohols react with halogenating agents:
Replacement of one halogen by another using halide salts:
Alkyl chloride/bromide with NaI in acetone:
\[
CH_3CH_2Br + NaI \xrightarrow{acetone} CH_3CH_2I + NaBr
\]
Used to prepare alkyl fluorides:
\[
RBr + AgF \rightarrow RF + AgBr
\]
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:
\[
C_2H_5COOAg + Br_2 \rightarrow C_2H_5Br + CO_2 + AgBr
\]
Replacement of halogen by a nucleophile (\( Nu^- \)):
SN1 Mechanism (Unimolecular):
SN2 Mechanism (Bimolecular):
When a haloalkane is heated with alcoholic KOH, it undergoes β-elimination to form an alkene.
\[
R-CH_2CHX-R' + alc.\ KOH \xrightarrow{\Delta} R-CH=CH-R' + KX + H_2O
\]
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)
Haloalkanes react with sodium metal in dry ether to form higher alkanes.
\[
2R-X + 2Na \xrightarrow{dry\ ether} R-R + 2NaX
\]
Example:
\[
2CH_3Br + 2Na \rightarrow C_2H_6 + 2NaBr
\]
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.
\[
CH_3CH_2OH + 4[O] \rightarrow CH_3COOH + 2H_2O
\]
\[
CH_3COCH_3 + 3Cl_2 + 4NaOH \rightarrow CHCl_3 + CH_3COONa + 3NaCl + 2H_2O
\]
\[
CH_3CH_2OH + I_2 + 4NaOH \rightarrow CHI_3 + HCOONa + 3NaI + 2H_2O
\]
Aryl diazonium salts react with Cu(I) halide to give haloarenes.
\[
C_6H_5N_2^+Cl^- + CuCl \rightarrow C_6H_5Cl + N_2 \uparrow
\]
\[
C_6H_5N_2^+Cl^- + CuBr \rightarrow C_6H_5Br + N_2 \uparrow
\]
Similar to Sandmeyer’s but uses Cu powder and halogen (Cl₂ or Br₂) directly:
\[
C_6H_5N_2^+Cl^- + Cu + Cl_2 \rightarrow C_6H_5Cl + N_2 + CuCl
\]
Halogenation of benzene in presence of Lewis acid (FeCl₃/AlCl₃):
\[
C_6H_6 + Cl_2 \xrightarrow{FeCl_3} C_6H_5Cl + HCl
\]
\[
C_6H_6 + Br_2 \xrightarrow{FeBr_3} C_6H_5Br + HBr
\]
Example (boiling point trend):
\[
CH_3Cl < CH_3Br < CH_3I
\]
Reason for water insolubility:
\[
\text{Energy released in forming new dipole-dipole interactions is less than energy needed to break hydrogen bonds in water}
\]
Directive Influence: Halogens are ortho/para-directing due to resonance, but deactivating due to –I effect.
\[
C_6H_5Cl + Cl_2 \xrightarrow{FeCl_3} o\text{-}ClC_6H_4Cl + p\text{-}ClC_6H_4Cl
\]
\[
C_6H_5Cl + HNO_3 \xrightarrow{H_2SO_4} o\text{-}NO_2C_6H_4Cl + p\text{-}NO_2C_6H_4Cl
\]
\[
C_6H_5Cl + H_2SO_4 \xrightarrow{heat} o\text{-}HSO_3C_6H_4Cl + p\text{-}HSO_3C_6H_4Cl
\]
Haloarenes are less reactive toward nucleophilic substitution due to delocalization of lone pairs on halogen.
\[
C_6H_5Cl + NaOH \xrightarrow{300^\circ C,\ 300\ atm} C_6H_5OH + NaCl
\]
\[
C_6H_5Cl + NH_3 \xrightarrow{Cu_2O,\ 200^\circ C} C_6H_5NH_2 + HCl
\]
Haloarenes can be reduced using hydrogen and catalyst:
\[
C_6H_5Cl + 2[H] \xrightarrow{H_2,\ Pd} C_6H_6 + HCl
\]
Reaction of haloarene and haloalkane with sodium in dry ether:
\[
C_6H_5Br + CH_3Br + 2Na \xrightarrow{dry\ ether} C_6H_5CH_3 + 2NaBr
\]
Two haloarenes + sodium in dry ether → biphenyl:
\[
2C_6H_5Br + 2Na \xrightarrow{dry\ ether} C_6H_5{-}C_6H_5 + 2NaBr
\]
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.
DDT = Dichlorodiphenyltrichloroethane
Structural formula:
\[
(ClC_6H_4)_2CHCCl_3
\]
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.
Using Lucas Test:
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:
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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.
Based on the number of alkyl groups attached to the carbon bearing the –OH group:
–OH group attached to a carbon bonded to one other carbon.
Example: Ethanol \( CH_3CH_2OH \)
CH₃−CH₂−OH
–OH group on a carbon bonded to two other carbon atoms.
Example: Isopropanol \( CH_3CHOHCH_3 \)
CH₃−CH(OH)−CH₃
–OH group on a carbon bonded to three carbon atoms.
Example: Tert-butanol \( (CH_3)_3COH \)
(CH₃)₃C−OH
\[
CH_2=CH_2 + H_2O \xrightarrow{H_2SO_4} CH_3CH_2OH
\]
(Ethene to ethanol using acid catalyst)
\[
CH_2=CH_2 \xrightarrow[NaBH_4]{Hg(OAc)_2,\ H_2O} CH_3CH_2OH
\]
\[
CH_2=CH_2 \xrightarrow[B_2H_6]{THF} CH_3CH_2BH_2 \xrightarrow[OH^-]{H_2O_2} CH_3CH_2OH
\]
\[
R-MgX + H_2C=O \xrightarrow{ether} RCH_2OMgX \xrightarrow{H_2O} RCH_2OH
\]
(Grignard + formaldehyde gives primary alcohol)
\[
R-X + OH^- \rightarrow R-OH + X^-
\]
Example:
\[
CH_3CH_2Br + NaOH \rightarrow CH_3CH_2OH + NaBr
\]
\[
RCHO + 2[H] \xrightarrow{LiAlH_4} RCH_2OH
\]
\[
RCOR' + 2[H] \rightarrow RCH(OH)R'
\]
By diazotization and hydrolysis:
\[
RNH_2 + HNO_2 \rightarrow ROH + N_2 + H_2O
\]
Bosch Process:
\[
CO + 2H_2 \xrightarrow{ZnO/Cr_2O_3,\ 300^\circ C,\ 300\ atm} CH_3OH
\]
From Glucose using Yeast:
\[
C_6H_{12}O_6 \xrightarrow{yeast} 2C_2H_5OH + 2CO_2
\]
Alcohols can donate a proton (H⁺) from the hydroxyl group, showing weak acidity:
\[
ROH + Na \rightarrow RONa + \frac{1}{2}H_2 \uparrow
\]
Alcohols react with active metals like sodium to form alkoxides:
\[
2CH_3CH_2OH + 2Na \rightarrow 2CH_3CH_2ONa + H_2 \uparrow
\]
Alcohols react with carboxylic acids in the presence of an acid catalyst (H₂SO₄) to form esters.
\[
RCOOH + R'OH \xrightarrow{H^+} RCOOR' + H_2O
\]
Alcohols react with hydrogen halides to form alkyl halides:
\[
R-OH + HX \rightarrow R-X + H_2O
\]
Example:
\[
CH_3CH_2OH + HBr \rightarrow CH_3CH_2Br + H_2O
\]
Alcohols react with acid chlorides to give esters:
\[
RCOCl + R'OH \rightarrow RCOOR' + HCl
\]
With acid anhydrides:
\[
(RCO)_2O + R'OH \rightarrow RCOOR' + RCOOH
\]
\[
CH_3CH_2OH \xrightarrow{[O]} CH_3CHO \xrightarrow{[O]} CH_3COOH
\]
Alcohols dehydrate in acidic medium to give alkenes.
\[
CH_3CH_2OH \xrightarrow{H_2SO_4,\ 170^\circ C} CH_2=CH_2 + H_2O
\]
Oxidation: Primary alcohol is oxidized to aldehyde, then further to carboxylic acid.
\[
\mathrm{CH_3CH_2OH} \xrightarrow[\text{mild}]{[O]} \mathrm{CH_3CHO} \xrightarrow[\text{strong}]{[O]} \mathrm{CH_3COOH}
\]
Reduction: Carboxylic acid or aldehyde reduced back to primary alcohol.
\[
\mathrm{CH_3COOH} \xrightarrow{[H]} \mathrm{CH_3CH_2OH}
\]
Oxidation of primary alcohol to aldehyde, then rearrangement to secondary alcohol via Grignard reagent.
Example: Using Grignard reagent with formaldehyde
\[
\mathrm{CH_3MgBr} + \mathrm{HCHO} \xrightarrow{H^+} \mathrm{CH_3CH_2OH}
\]
Secondary alcohol can be converted to tertiary alcohol by reaction with Grignard reagent:
\[
\mathrm{CH_3CHOHCH_3} + \mathrm{CH_3MgBr} \xrightarrow{H^+} (CH_3)_3COH
\]
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 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.
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:
The reaction proceeds by the formation of a carbocation intermediate. The stability of the carbocation increases from primary to tertiary, explaining the reaction speed:
| 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) |
Nomenclature: The –OH group attached to benzene ring forms the suffix phenol or prefix hydroxy- if another principal group exists.
Phenol is prepared by the hydrolysis of benzene diazonium chloride:
Note: The diazonium salt is prepared by treating aniline with nitrous acid at 0–5°C.
Chlorobenzene is fused with sodium hydroxide at high temperature and pressure to give sodium phenoxide, which on acidification yields phenol:
Benzene sulphonic acid on fusion with sodium hydroxide produces sodium phenoxide, which upon acidification gives phenol:
Phenol is industrially manufactured by the cumene process, which involves the following steps:
Benzene reacts with propylene in presence of an acid catalyst (like AlCl₃) to form cumene:
Cumene is oxidized by air to form cumene hydroperoxide:
Cumene hydroperoxide is cleaved in acidic medium to give phenol and acetone:
Phenol is a white crystalline solid at room temperature with a pleasant odor.
The hydroxyl group (\( -OH \)) in phenol forms hydrogen bonds, which contributes to its solubility in polar solvents.
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:
Phenol reacts with sodium hydroxide to form sodium phenoxide and water:
Phenol reacts with sodium metal to give sodium phenoxide and hydrogen gas:
When heated with zinc dust, phenol forms benzene:
Phenol reacts with acetyl chloride or acetic anhydride to form esters (acetates):
Phenol reacts with PCl₅ to form phenyl chloride and phosphorus oxychloride:
Phenol undergoes electrophilic substitution reactions readily due to activation by the –OH group:
Phenol reacts with carbon dioxide in presence of sodium hydroxide under pressure and heat to form sodium salicylate, which on acidification gives salicylic acid:
Phenol reacts with chloroform in the presence of sodium hydroxide to form salicylaldehyde (o-hydroxybenzaldehyde):
Aliphatic ethers have the general formula: \( R{-}O{-}R' \), where \( R \) and \( R' \) are alkyl groups.
Structure consists of an oxygen atom bonded to two alkyl groups.
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.
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.
Prepared by Williamson Ether Synthesis:
\( C_6H_5O^- + CH_3I \rightarrow C_6H_5OCH_3 + I^- \)
(Phenoxide ion reacts with methyl iodide)
The ether oxygen activates the aromatic ring toward electrophilic substitution by donating electron density via resonance.
Both aldehydes and ketones contain the carbonyl group \( >C=O \), where the carbon is sp² hybridized and planar.
Nucleophile attacks the electrophilic carbon of the carbonyl group, pushing electrons onto oxygen.
The alpha hydrogen (adjacent to the carbonyl group) is acidic due to resonance stabilization of the enolate ion:
This property enables aldol condensation and related reactions.
- Primary alcohols are oxidized to aldehydes.
- Secondary alcohols are oxidized to ketones.
Cleavage of double bonds by ozone followed by hydrolysis gives aldehydes/ketones.
Terminal alkynes give aldehydes; internal alkynes give ketones.
Rosenmund’s Reduction: Acid chlorides to aldehydes using hydrogen over Pd/BaSO₄.
Reaction with Dialkyl Cadmium: Acid chlorides react with dialkyl cadmium to form ketones.
Dry distillation of calcium salts of fatty acids gives ketones.
Stephen Reaction: Nitriles are reduced to imines, then hydrolyzed to aldehydes.
Grignard’s Reagent: Nitriles react with RMgX and then are hydrolyzed to ketones.
Grignard reagent reacts with esters to form ketones (intermediate stage before tertiary alcohol).
Oxidation with Chromyl Chloride (Etard Reaction):
This reaction is performed in a dry solvent like carbon tetrachloride, followed by hydrolysis to yield benzaldehyde.
In presence of CN⁻, two molecules of benzaldehyde condense to form benzoin.
Reaction with acetic anhydride and sodium acetate to form cinnamic acid.
General formula: \( R{-}COOH \)
The carboxyl group consists of a carbon double bonded to an oxygen and single bonded to a hydroxyl group.
Carboxylic acids are weak acids due to resonance stabilization of the carboxylate ion:
Acidity is affected by the inductive effect and hydrogen bonding.
Carboxylic acids are weak acids. The presence of electron-withdrawing groups increases their acidity, while electron-donating groups decrease it.
Carboxylic acids react with alcohols to form esters in a reaction known as esterification.
Carboxylic acids can lose a CO₂ molecule when heated with soda lime (NaOH + CaO) to form alkanes.
Carboxylic acids undergo α-halogenation in the presence of halogens (Cl₂ or Br₂) and red phosphorus.
The carboxyl group is meta-directing and deactivates the benzene ring towards electrophilic substitution.
General formula: \( R{-}NH_2 \) (for primary amines)
The nitrogen atom in amines is \( sp^3 \)-hybridized with a lone pair, giving a trigonal pyramidal geometry.
Amines contain a nitrogen atom with a lone pair, making them nucleophilic.
General formula (for aliphatic primary amines): \( \ce{C_nH_{2n+1}NH_2} \)
Reduction of alcohol to amine (via alkyl halide or nitrile):
Alkyl halide reacts with excess ammonia (ammonolysis):
Reduction of nitrile using hydrogen or lithium aluminium hydride:
Amide reacts with bromine and alkali to form a primary amine with one carbon less:
Reduction of nitroalkane or nitrobenzene to amine:
Used to prepare primary amines (especially aliphatic):
Basicity is due to lone pair of electrons on nitrogen.
Order in aqueous solution:
Amines react with alkyl halides to form higher amines and quaternary salts:
Excess alkyl halide can lead to quaternary ammonium salt:
1° and 2° amines react with acid chlorides or anhydrides:
Mechanism involves nucleophilic attack of lone pair on carbonyl carbon.
Used to test for 1° amines (foul-smelling isocyanide formed):
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By reduction of nitrobenzene:
Aniline forms salts with mineral acids like HCl and H2SO4:
Reacts with acetic anhydride to form acetanilide (less reactive amide):
Reaction with alkyl halides to form alkylated anilines:
Reaction with benzoyl chloride:
Formation of isocyanide (test for primary amines):
Reaction with nitrous acid to form diazonium salt:
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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.)
Used to replace the diazonium group with halogens.
\[
C_6H_5N_2^+Cl^- + CuCl \rightarrow C_6H_5Cl + N_2 \uparrow
\]
Similar to Sandmeyer but uses Cu powder and HCl.
\[
C_6H_5N_2^+Cl^- + Cu + HCl \rightarrow C_6H_5Cl + N_2 + CuCl
\]
Used to replace the diazonium group with fluorine.
\[
C_6H_5N_2^+BF_4^- \xrightarrow{\Delta} C_6H_5F + BF_3 + N_2
\]
Diazonium salts undergo electrophilic substitution with phenols and aromatic amines to form azo compounds.
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Carbohydrates are organic compounds composed of carbon, hydrogen, and oxygen in the ratio Cn(H2O)n. They are also called saccharides.
D and L refer to the configuration around the asymmetric carbon farthest from the carbonyl group, based on glyceraldehyde.
Open-chain Glucose:
\[ \ce{HOCH2(CHOH)4CHO} \]
Cyclic α-D-glucose:
\[ \ce{C6H12O6} \] forms a six-membered ring (pyranose).
Fructose (Open-chain):
\[ \ce{CH2OH-CO-(CHOH)3-CH2OH} \]
Bromine Water Test (for Glucose):
\[ \ce{C6H12O6 + Br2 + H2O -> C6H12O7 + 2HBr} \]
Fructose does not react directly with bromine water.
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Proteins are polymers made up of amino acids linked by peptide bonds forming polypeptides.
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).
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Biological catalysts that speed up chemical reactions without being consumed.
Mostly proteins with specific active sites for substrate binding.
Chemical messengers that regulate physiology and behavior.
Examples include insulin (protein hormone) and steroid hormones.
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Nucleic acids are polymers made of nucleotides. Each nucleotide consists of:
DNA is a double helical structure discovered by Watson and Crick.
It consists of two strands running in opposite directions (antiparallel) held together by hydrogen bonds between complementary bases:
Backbone is made of sugar (deoxyribose) and phosphate groups.
Formula example for deoxyribose sugar: \\[ \ce{C5H10O4} \\]
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.
| 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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For decades, the PIN code (Postal Index Number) has been an integral part of India’s postal system, helping identify broad geographical regions for mail delivery. But now, the Indian Postal Department is ushering in a new era of precision with the launch of DigiPIN, a 10-digit digital address system designed to pinpoint your exact location — whether it’s your home, clinic, office, or business. What is DigiPIN? DigiPIN (Digital PIN) is a geo-coded digital address system introduced by India Post, in collaboration with IIT Hyderabad, NRSC, and ISRO. Unlike traditional PIN codes that cover entire zones or localities, DigiPIN identifies your location within a 4-meterRead More →
1.Electrostatics (i) Electric Charges and Fields Electric Charges – ISC Physics Class 12 Electric Charges Definition: Electric charge is a fundamental property of matter carried by some elementary particles that causes them to experience a force when placed in an electromagnetic field. Types of Charges Positive Charge: Deficiency of electrons (e.g., proton). Negative Charge: Excess of electrons (e.g., electron). Basic Properties of Electric Charge Additivity of Charges: Total charge is the algebraic sum of individual charges. Quantization of Charge: Charge (q) is always an integral multiple of the elementary charge (e). q = ±ne where n is an integer and e = 1.6 × 10⁻¹⁹ CRead More →