CBSE Class 12 Chemistry Part 1 Exercise Solutions

How to Use This Page

This page provides structured, exam-ready answers for all four chapters in CBSE Class 12 Chemistry Part 1. Each section mirrors the question patterns you will encounter in the board exam: MCQs, fill-in-the-blanks, true/false, match the following, short answers, and detailed answers with worked numericals.

Before you begin, read the relevant NCERT chapter thoroughly. Then attempt the questions on your own. Only after you have written your own answer should you compare it with the solutions given here. Pay special attention to the keywords, formulas, and step-by-step methods shown in the detailed answers — examiners award marks for correct methodology and key terms, not for memorised reproductions.

For numerical problems, always write the given data, the formula used, the substitution step, and the final answer with correct units and significant figures. A numerical answer without units or with wrong significant figures will lose marks even if the calculation is correct.

Use the Key Concepts section as a rapid revision tool before exams. The Common Mistakes and Exam Checklist sections will help you avoid the errors that cost the most marks.

Key Concepts (Quick Recap)

Chapter 1: Solutions

• Types of solutions: solid, liquid, gaseous — classified by the physical state of solute and solvent
• Molarity (M) = moles of solute / volume of solution in litres
• Molality (m) = moles of solute / mass of solvent in kg
• Mole fraction (x) = moles of component / total moles; x\_A + x\_B = 1
• Raoult's law for volatile solutes: p\_A = x\_A \* p\_A^0 (partial pressure = mole fraction x vapour pressure of pure component)
• Raoult's law for non-volatile solutes: Relative lowering of vapour pressure, (p\_A^0 - p\_A) / p\_A^0 = x\_B = n\_B / (n\_A + n\_B)
• Ideal solutions: Obey Raoult's law over entire composition range; delta H\_mix = 0, delta V\_mix = 0
• Non-ideal solutions — positive deviation: p\_A > x\_A p\_A^0; delta H\_mix > 0 (e.g., ethanol + water)
• Non-ideal solutions — negative deviation: p\_A < x\_A p\_A^0; delta H\_mix < 0 (e.g., chloroform + acetone)
• Colligative properties (depend on number of solute particles, not their nature):
• Relative lowering of vapour pressure: (p\_A^0 - p\_A) / p\_A^0 = n\_B / n\_A = (W\_B / M\_B) / (W\_A / M\_A)
• Elevation of boiling point: delta T\_b = K\_b \* m; where K\_b = ebullioscopic constant
• Depression of freezing point: delta T\_f = K\_f \* m; where K\_f = cryoscopic constant
• Osmotic pressure: pi = nRT / V = CRT; where R = 0.0821 L atm K^-1 mol^-1
• Van't Hoff factor (i): i = observed colligative property / calculated colligative property = (normal molar mass) / (observed molar mass) = (total particles after dissociation) / (total particles before dissociation)
• For dissociation: i = 1 + (alpha)(n - 1), where alpha = degree of dissociation, n = number of ions
• For association: i = 1 - (1 - 1/n)alpha
• Abnormal molar mass: M\_observed = M\_normal / i
• Henry's law: p = K\_H \* x (solubility of a gas is proportional to the pressure of the gas above the solution)
• Osmosis: spontaneous flow of solvent from low solute concentration to high solute concentration through a semipermeable membrane
• Isotonic solutions: two solutions with the same osmotic pressure

Chapter 2: Electrochemistry

• Conductivity (kappa) = 1 / rho; unit: S m^-1 or S cm^-1
• Molar conductivity (Lambda\_m) = kappa / c; unit: S cm^2 mol^-1
• Kohlrausch's law: Lambda\_m^0 = v\_+ \ lambda\_+^0 + v\_- \ lambda\_-^0 (limiting molar conductivity = sum of ionic contributions)
• Debye-Huckel-Onsager equation: Lambda\_m = Lambda\_m^0 - A \* sqrt(c) (for strong electrolytes)
• Weak electrolytes: Lambda\_m increases steeply with dilution because degree of dissociation increases
• Strong electrolytes: Lambda\_m increases slightly with dilution; approaches Lambda\_m^0 at infinite dilution
• Nernst equation: E\_cell = E\_cell^0 - (RT / nF) \ ln Q = E\_cell^0 - (0.0591 / n) \ log Q at 298 K
• Standard electrode potential: E^0 values from standard tables; more positive E^0 = stronger oxidising agent
• E\_cell^0 = E\_cathode^0 - E\_anode^0
• Gibbs free energy: delta G^0 = -nF E\_cell^0; also delta G^0 = -2.303 RT \* log K\_c
• Relationship: E\_cell^0 = (0.0591 / n) \* log K\_c at 298 K
• Faraday's first law: m = Z \ I \ t = (M / nF) \ I \ t (mass deposited = electrochemical equivalent x current x time)
• Faraday's second law: masses deposited by same quantity of electricity are in the ratio of their chemical equivalents (equivalent weights)
• 1 Faraday (F) = 96485 C mol^-1 (approx 96500 C mol^-1)
• Resistivity (rho) = R \* A / l; unit: ohm cm or ohm m
• Conductance (G) = 1 / R; unit: S (siemens) or mho
• *Cell constant (G\)** = l / A = kappa / G; determined using a standard KCl solution

Chapter 3: Chemical Kinetics

• Rate of reaction: r = -(1/a)(d[A]/dt) = -(1/b)(d[B]/dt) = (1/c)(d[C]/dt) = (1/d)(d[D]/dt) for aA + bB → cC + dD
• Rate law: Rate = k[A]^m[B]^n; where m, n are orders determined experimentally (NOT from stoichiometry)
• Order of reaction: sum of powers of concentration terms in the rate law; can be 0, 1, 2, fractional
• Molecularity: number of reactant molecules participating in an elementary step; always a whole number (1, 2, or 3)
• Zero order: Rate = k; [A] = [A]\_0 - kt; t\_1/2 = [A]\_0 / 2k
• First order: Rate = k[A]; ln[A] = ln[A]\_0 - kt; k = (2.303 / t) \* log([A]\_0 / [A]); t\_1/2 = 0.693 / k
• Second order (equal concentrations): Rate = k[A]^2; 1/[A] = 1/[A]\_0 + kt; t\_1/2 = 1 / (k[A]\_0)
• Pseudo first order: When one reactant is in large excess, its concentration remains nearly constant (e.g., hydrolysis of ester in aqueous solution)
• Arrhenius equation: k = A \* e^(-Ea / RT); ln k = ln A - Ea / RT
• Arrhenius (two-temperature form): ln(k\_2 / k\_1) = (Ea / R) \* (1/T\_1 - 1/T\_2)
• Activation energy (Ea): minimum energy required for a reaction to occur; determined from slope of ln k vs 1/T plot (slope = -Ea/R)
• Collision theory: Rate = P \ Z \ e^(-Ea / RT); where Z = collision frequency, P = steric factor
• Effect of temperature: rate approximately doubles for every 10 K rise in temperature (rough rule)
• Units of k: zero order — mol L^-1 s^-1; first order — s^-1; second order — L mol^-1 s^-1; nth order — (mol L^-1)^(1-n) s^-1

Chapter 4: Surface Chemistry

• Adsorption: accumulation of molecular species at the surface (not inside the bulk)
• Absorption: uniform penetration of molecules throughout the bulk
• Physisorption: weak van der Waals forces, reversible, multimolecular, decreases with temperature, no activation energy
• Chemisorption: strong chemical bonds, irreversible (often), unimolecular, increases then decreases with temperature, requires activation energy
• Freundlich adsorption isotherm: x/m = k \ p^(1/n) (for gases); x/m = k \ C^(1/n) (for solutions); log(x/m) = log k + (1/n) \* log p
• Langmuir adsorption isotherm: theta = (b \ p) / (1 + b \ p); based on monolayer adsorption on homogeneous surface
• Catalysis: process of altering the rate of a reaction by adding a catalyst
• Homogeneous catalyst: same phase as reactants (e.g., NO in Lead Chamber process)
• Heterogeneous catalyst: different phase (e.g., V₂O₅ in Contact process, Fe in Haber's process)
• Enzymes: biological catalysts; protein in nature; specific, temperature-sensitive (optimum ~ 37°C), pH-sensitive
• Colloids: heterogeneous system where one phase (dispersed phase) is dispersed in another (dispersion medium)
• Lyophilic colloids: solvent-attracting (e.g., gum, gelatin in water); reversible, stable, coagulated only by high electrolyte concentration
• Lyophobic colloids: solvent-repelling (e.g., metals in water); irreversible, unstable, coagulated by small amounts of electrolyte
• Tyndall effect: scattering of light by colloidal particles (distinguishes colloidal from true solution)
• Brownian motion: random zigzag motion of colloidal particles due to unbalanced bombardment by dispersion medium molecules
• Electrophoresis: movement of colloidal particles under an applied electric field (determines charge)
• Coagulation: precipitation of a colloid by adding electrolyte; minimum concentration of electrolyte needed = coagulation value
• Hard Schulze rule: greater the valence of the flocculating ion, greater is the coagulating power; Al³⁺ > Ba²⁺ > Na⁺ for negatively charged sols
• Emulsion: colloid of two immiscible liquids; oil-in-water (O/W) and water-in-oil (W/O); stabilised by emulsifying agents (e.g., soap, detergent)
• Gold number: measure of protective power of a lyophilic colloid; smaller gold number = greater protective power
• Peptization: converting a freshly precipitated substance into a colloidal sol by adding an electrolyte (peptizing agent)

I. Choose the Best Answer (MCQs)

1. Which of the following is a colligative property?

(a) Boiling point (b) Osmotic pressure (c) Surface tension (d) Vapour pressure of pure solvent

2. The van't Hoff factor for 0.1 M BaCl₂ solution assuming complete dissociation is:

(a) 0.33 (b) 1 (c) 2 (d) 3

3. The molar conductivity of a weak electrolyte at infinite dilution can be determined by:

(a) Extrapolation of Lambda\_m vs sqrt(c) plot (b) Kohlrausch's law using ionic conductivities (c) Direct measurement at very low concentration (d) Using Nernst equation

4. For a first-order reaction, the half-life:

(a) Depends on initial concentration (b) Is inversely proportional to initial concentration (c) Is independent of initial concentration (d) Is proportional to the rate constant

5. Which of the following is an example of chemisorption?

(a) H₂ gas on charcoal (b) N₂ gas on iron surface (c) O₂ gas on silica gel (d) H₂O vapour on activated charcoal

6. The standard electrode potential of Cu²⁺/Cu is +0.34 V and that of Zn²⁺/Zn is -0.76 V. The EMF of a Daniel cell is:

(a) 0.42 V (b) 1.10 V (c) 0.76 V (d) 0.34 V

7. In the Freundlich adsorption isotherm, x/m = k \* p^(1/n), the value of 1/n at high pressure is:

(a) 1 (b) 0 (c) Between 0 and 1 (d) Greater than 1

8. The correct relationship between Gibbs free energy and cell EMF is:

(a) delta G = +nFE (b) delta G = -nFE (c) delta G = nF/E (d) delta G = -E/nF

9. Which has the highest coagulating power for a negatively charged sol?

(a) NaCl (b) BaCl₂ (c) AlCl₃ (d) K₄[Fe(CN)₆]

10. For a zero-order reaction A → products, the rate constant k has the units:

(a) s⁻¹ (b) L mol⁻¹ s⁻¹ (c) mol L⁻¹ s⁻¹ (d) mol⁻¹ L s⁻¹

II. Fill in the Blanks

1. The osmotic pressure of a solution is related to its molar concentration by the equation pi = CRT, where R is the gas constant and T is the temperature in kelvin.

2. In a galvanic cell, oxidation takes place at the anode and reduction takes place at the cathode.

3. The standard electrode potential of the hydrogen electrode is assigned a value of 0 V (zero volt).

4. The half-life of a zero-order reaction is directly proportional to the initial concentration of the reactant.

5. According to Kohlrausch's law, the molar conductivity of an electrolyte at infinite dilution is equal to the sum of the ionic conductivities of its constituent ions.

6. The Tyndall effect is caused by scattering of light by colloidal particles.

7. In an oil-in-water (O/W) emulsion, the dispersed phase is oil and the dispersion medium is water.

8. The Arrhenius equation can be written as k = A e^(-Ea/RT), where A is the pre-exponential factor (frequency factor) and Ea is the activation energy.

9. The degree of dissociation alpha of a weak electrolyte can be calculated from molar conductivities as alpha = Lambda\_m / Lambda\_m^0.

10. The protective power of a lyophilic colloid is measured in terms of its gold number; the smaller the gold number, the greater is the protective power.

III. True or False

1. True or False: In an ideal solution, the enthalpy of mixing is zero.

True. In an ideal solution, A-A, B-B, and A-B intermolecular forces are nearly equal, so delta H\_mix = 0 and delta V\_mix = 0.

2. True or False: The conductivity of a solution decreases with dilution.

True. Conductivity (kappa) decreases with dilution because the number of ions per unit volume decreases. However, molar conductivity (Lambda\_m) increases with dilution because interionic attractions decrease.

3. True or False: The order of a reaction can be determined from the stoichiometric coefficients of the balanced equation.

False. The order of a reaction is determined experimentally from the rate law; it cannot be deduced from stoichiometric coefficients. Only for an elementary (single-step) reaction does the order coincide with stoichiometry.

4. True or False: Physisorption is always exothermic and increases with increasing temperature.

False. Physisorption is exothermic but decreases with increasing temperature because it involves weak van der Waals forces that are disrupted by thermal energy. Chemisorption initially increases with temperature (due to activation energy requirement) and then decreases.

5. True or False: The van't Hoff factor for an associating solute is greater than 1.

False. For an associating solute (e.g., acetic acid in benzene forms dimers), the van't Hoff factor i is less than 1 because the effective number of particles decreases. For dissociating solutes (e.g., NaCl in water), i > 1.

6. True or False: In a concentration cell, the two half-cells contain the same electrolyte at different concentrations.

True. A concentration cell generates EMF from the concentration difference between two identical half-cells. The electrode in the dilute solution acts as the anode (oxidation) and the electrode in the concentrated solution acts as the cathode (reduction).

7. True or False: Brownian motion in colloidal particles is caused by unbalanced bombardment by molecules of the dispersion medium.

True. The random, zigzag Brownian motion arises because colloidal particles are simultaneously hit by a large number of medium molecules from different directions, and the forces do not perfectly cancel at any instant.

8. True or False: The rate constant of a first-order reaction has units of mol L⁻¹ s⁻¹.

False. The rate constant of a first-order reaction has units of s⁻¹ (or min⁻¹, yr⁻¹ etc.) because the rate law is Rate = k[A], so k = Rate / [A] = (mol L⁻¹ s⁻¹) / (mol L⁻¹) = s⁻¹. The units mol L⁻¹ s⁻¹ belong to a zero-order rate constant.

IV. Match the Following

Key:

• 1 → (e) Raoult's law relates the vapour pressure of a solution component to its mole fraction.
• 2 → (f) Kohlrausch's law states that limiting molar conductivity is the sum of ionic conductivities.
• 3 → (a) The Arrhenius equation relates the rate constant to activation energy and temperature.
• 4 → (c) The Hardy-Schulze rule states that coagulating power increases with valence of the flocculating ion.
• 5 → (b) The Nernst equation calculates electrode potential under non-standard conditions of concentration and temperature.
• 6 → (d) The Freundlich isotherm describes adsorption as x/m = k p^(1/n).

V. Answer Briefly

Q1. Define an ideal solution. What are the conditions for a solution to be ideal?

An ideal solution obeys Raoult's law over the entire range of composition and at all temperatures. The conditions are: (i) The intermolecular forces between A-A, B-B, and A-B are nearly identical, so delta H\_mix = 0. (ii) The volumes are additive, so delta V\_mix = 0. Examples include benzene + toluene and hexane + heptane.

Q2. What is the difference between order and molecularity of a reaction?

| Feature | Order | Molecularity | |---|---|---| | Definition | Sum of powers of concentration terms in experimentally determined rate law | Number of reactant molecules participating in an elementary step | | Determination | Experimental | Theoretical (from reaction mechanism) | | Value | Can be zero, fractional, or negative | Always a positive integer (1, 2, or 3) | | Applicability | Valid for complex (multi-step) reactions | Meaningful only for elementary reactions |

Q3. Why does the conductivity of a solution decrease with dilution, but molar conductivity increases?

Conductivity (kappa) is the conductance of 1 cm³ of the solution. On dilution, the number of ions per unit volume decreases, so conductivity decreases. Molar conductivity (Lambda\_m = kappa / c) increases because: (i) degree of dissociation increases (more ions per mole), and (ii) interionic attractions decrease (ions move more freely). The decrease in kappa is less than proportional to the decrease in concentration, so Lambda\_m rises.

Q4. What is a buffer solution? How does it resist change in pH?

A buffer solution resists change in pH upon addition of small amounts of acid or base. It consists of a weak acid and its conjugate base (e.g., CH₃COOH + CH₃COONa), or a weak base and its conjugate acid (e.g., NH₄OH + NH₄Cl). When acid is added, the conjugate base neutralises it; when base is added, the weak acid neutralises it. The pH remains nearly constant because the ratio [acid]/[salt] (or [base]/[salt]) changes very little.

Q5. Differentiate between physisorption and chemisorption.

| Property | Physisorption | Chemisorption | |---|---|---| | Forces | Weak van der Waals | Strong chemical bonds | | Enthalpy | 20-40 kJ/mol | 80-240 kJ/mol | | Reversibility | Reversible | Irreversible (often) | | Temperature effect | Decreases with T | Increases then decreases | | Layer | Multimolecular | Unimolecular | | Specificity | Non-specific | Highly specific | | Activation energy | Negligible | Significant |

Q6. What is the Tyndall effect? How does it help distinguish a colloidal solution from a true solution?

The Tyndall effect is the scattering of a beam of light by colloidal particles, making the path of light visible when viewed from the side. This occurs because colloidal particles (1-1000 nm) are large enough to scatter visible light, while true solution particles (< 1 nm) are too small. To distinguish, shine a beam of light through the solution: in a true solution, the beam is invisible from the side; in a colloidal solution, the beam becomes visible as a bright cone (Tyndall cone).

Q7. State Faraday's laws of electrolysis.

First law: The mass of a substance deposited or liberated at an electrode during electrolysis is directly proportional to the quantity of electricity passed through the electrolyte. m = Z \* Q, where Z is the electrochemical equivalent. Second law: When the same quantity of electricity is passed through different electrolytes in series, the masses of substances deposited are proportional to their chemical equivalent weights (equivalent weight = molar mass / n, where n is the number of electrons transferred).

Q8. What is meant by the "abnormal molar mass" of a solute? When does it occur?

An abnormal molar mass is one that differs from the expected (theoretical) molar mass, as calculated from colligative properties. It occurs when the solute dissociates (giving i > 1, observed molar mass lower than expected, e.g., NaCl in water) or associates (giving i < 1, observed molar mass higher than expected, e.g., acetic acid in benzene). The van't Hoff factor i = M\_normal / M\_observed quantifies this deviation.

VI. Answer in Detail

Q1. Derive the expression for the elevation of boiling point and show how it is used to determine the molar mass of a solute. Calculate the boiling point elevation when 1.8 g of glucose (M = 180 g mol⁻¹) is dissolved in 100 g of water. (K\_b for water = 0.52 K kg mol⁻¹)

Derivation: When a non-volatile solute is added to a solvent, the vapour pressure of the solution decreases. The solution boils at a higher temperature than the pure solvent because a higher temperature is needed to make the vapour pressure equal to atmospheric pressure. The elevation of boiling point, delta T\_b = T\_b - T\_b^0, is proportional to the molality of the solution: delta T\_b = K\_b \ m where K\_b = ebullioscopic constant (molal elevation constant) and m = molality = (n\_B \ 1000) / (W\_A in grams). Since m = (W\_B / M\_B) \ (1000 / W\_A), we get: M\_B = (K\_b \ W\_B \ 1000) / (delta T\_b \ W\_A) This is how molar mass of an unknown solute can be determined experimentally. Calculation: Given: W\_B = 1.8 g, M\_B = 180 g mol⁻¹, W\_A = 100 g, K\_b = 0.52 K kg mol⁻¹ Step 1: Calculate molality n\_B = W\_B / M\_B = 1.8 / 180 = 0.01 mol m = n\_B / (W\_A / 1000) = 0.01 / (100 / 1000) = 0.01 / 0.1 = 0.1 mol kg⁻¹ Step 2: Calculate delta T\_b delta T\_b = K\_b \ m = 0.52 \ 0.1 = 0.052 K Step 3: Boiling point of solution T\_b = T\_b^0 + delta T\_b = 100 + 0.052 = 100.052°C The boiling point of the glucose solution is 100.052°C, an elevation of 0.052 K above pure water.

Q2. Calculate the EMF of the following cell at 298 K and determine whether the cell reaction is spontaneous:

Step 1: Identify the cell The cell notation shows Zn as the anode (oxidation) and Cu as the cathode (reduction). Anode (oxidation): Zn(s) → Zn²⁺(0.1 M) + 2e⁻ Cathode (reduction): Cu²⁺(0.01 M) + 2e⁻ → Cu(s) Overall: Zn(s) + Cu²⁺(0.01 M) → Zn²⁺(0.1 M) + Cu(s) Step 2: Calculate E⁰\_cell E⁰\_cell = E⁰\_cathode - E⁰\_anode = E⁰\_(Cu²⁺/Cu) - E⁰\_(Zn²⁺/Zn) = 0.34 - (-0.76) = 1.10 V Step 3: Apply the Nernst equation E\_cell = E⁰\_cell - (0.0591 / n) \ log(Q) where n = 2 (number of electrons transferred) Q = [Zn²⁺] / [Cu²⁺] = 0.1 / 0.01 = 10 E\_cell = 1.10 - (0.0591 / 2) \ log(10) E\_cell = 1.10 - (0.0591 / 2) \* 1 E\_cell = 1.10 - 0.02955 E\_cell = 1.07 V (approximately) Step 4: Determine spontaneity Since E\_cell = 1.07 V > 0, the cell reaction is spontaneous in the forward direction as written. Also, delta G = -nFE\_cell = -2 × 96500 × 1.07 = -206,510 J = -206.5 kJ, confirming spontaneity.

Q3. A first-order reaction has a rate constant of 2.0 × 10⁻³ s⁻¹ at 300 K and 4.0 × 10⁻³ s⁻¹ at 320 K. Calculate: (a) the half-life at 300 K, (b) the activation energy (Ea), and (c) the time required for 75% completion at 300 K.

Q4. Explain the colligative properties of solutions with the van't Hoff factor. A solution containing 1.0 g of KCl (M = 74.5 g mol⁻¹) in 100 g of water freezes at -0.48°C. Calculate the van't Hoff factor and the degree of dissociation of KCl. (K\_f for water = 1.86 K kg mol⁻¹)

Theory: Colligative properties depend on the number of solute particles, not their nature. For electrolytes that dissociate, the observed colligative property is greater than the calculated value because dissociation increases the number of particles. The van't Hoff factor i accounts for this: i = observed colligative property / calculated colligative property (assuming no dissociation) i = M\_normal / M\_observed For dissociation: i = 1 + alpha(n - 1), where alpha = degree of dissociation, n = number of ions produced Calculation: Given: W\_B = 1.0 g, M\_B = 74.5 g mol⁻¹, W\_A = 100 g = 0.1 kg, delta T\_f = 0.48°C (since pure water freezes at 0°C), K\_f = 1.86 K kg mol⁻¹ Step 1: Calculate the expected (theoretical) freezing point depression n\_B = W\_B / M\_B = 1.0 / 74.5 = 0.01342 mol m = n\_B / (W\_A in kg) = 0.01342 / 0.1 = 0.1342 mol kg⁻¹ delta T\_f(calculated) = K\_f × m = 1.86 × 0.1342 = 0.2496 K Step 2: Calculate the van't Hoff factor i = delta T\_f(observed) / delta T\_f(calculated) = 0.48 / 0.2496 = 1.923 Step 3: Calculate the degree of dissociation KCl → K⁺ + Cl⁻ (n = 2 ions) i = 1 + alpha(n - 1) 1.923 = 1 + alpha(2 - 1) 1.923 = 1 + alpha alpha = 1.923 - 1 = 0.923 or 92.3% KCl is approximately 92.3% dissociated in this solution.

Q5. What are emulsions? How are they classified? Explain the role of emulsifying agents with examples. Describe any three methods to identify the type of emulsion.

Emulsions are colloidal systems in which both the dispersed phase and the dispersion medium are liquids that are ordinarily immiscible. An emulsion is a liquid-in-liquid colloid. Classification: 1. Oil-in-water (O/W) emulsion: Oil droplets dispersed in water. Example: Milk (fat droplets in water), vanishing cream. 2. Water-in-oil (W/O) emulsion: Water droplets dispersed in oil. Example: Butter, cold cream, cod liver oil emulsion. Role of emulsifying agents: Emulsifying agents (or emulsifiers) are substances that stabilise emulsions by forming a protective film around the dispersed droplets, preventing them from coalescing. They reduce the interfacial tension between the two immiscible liquids. Examples: - Soap (sodium stearate): stabilises O/W emulsions; the hydrophilic end faces water, hydrophobic end faces oil. - Detergents: similar to soap, work as emulsifiers in O/W systems. - Gum arabic, gelatin: natural emulsifiers used in food products. - Tweens (polysorbates) and Spans (sorbitan esters): used in pharmaceutical emulsions. Methods to identify the type of emulsion: 1. Dye test: Add a water-soluble dye (e.g., methylene blue) to the emulsion. If the entire emulsion appears coloured uniformly, it is O/W (continuous phase is water, which takes up the dye). If only the dispersed droplets appear coloured under a microscope, it is W/O. 2. Dilution test: Add water to the emulsion. If it mixes easily and dilutes, the continuous phase is water → O/W emulsion. If water does not mix in, add oil; easy mixing indicates W/O emulsion. 3. Conductivity test: An O/W emulsion conducts electricity because water (continuous phase) is a conductor. A W/O emulsion shows very low conductivity because oil is the continuous phase and is a poor conductor. 4. Filter paper test: Place a drop on filter paper. O/W emulsion wets the paper (water spreads); W/O emulsion leaves an oily stain.

Q6. Explain the construction, working, and applications of a standard hydrogen electrode (SHE). Why is it needed in electrochemistry?

Construction: The standard hydrogen electrode consists of a platinum electrode coated with platinum black (to increase surface area), dipped in a 1 M HCl solution (standard H⁺ ion concentration), with hydrogen gas at 1 atm pressure bubbling over it at 298 K. The platinum electrode is inert — it does not participate in the reaction but provides a surface for the redox process. Representation: Pt | H₂(g, 1 atm) | H⁺(aq, 1 M) Working: When connected to another half-cell, the hydrogen electrode acts as either anode or cathode depending on the other electrode's potential: - As anode (oxidation): H₂(g) → 2H⁺(aq) + 2e⁻ (if the other electrode has a higher reduction potential) - As cathode (reduction): 2H⁺(aq) + 2e⁻ → H₂(g) (if the other electrode has a lower reduction potential) Standard electrode potential of SHE: By international convention, E⁰ for the standard hydrogen electrode is assigned a value of 0 V at all temperatures. All other standard electrode potentials are measured relative to SHE. Why SHE is needed: 1. It provides a universal reference point. Since absolute electrode potentials cannot be measured, all electrode potentials are measured relative to SHE. 2. It allows comparison of the oxidising and reducing tendencies of different substances on a common scale. 3. It enables the calculation of EMF of any galvanic cell: E\_cell = E\_cathode - E\_anode, where both values are relative to SHE. Applications: 1. Determination of standard electrode potentials of other half-cells. 2. Measurement of pH of solutions (using the hydrogen electrode as a pH electrode; E = -0.0591 × pH at 298 K). 3. As a reference electrode in potentiometric titrations. Limitations: 1. Requires pure hydrogen gas at 1 atm — difficult to maintain. 2. Platinum surface can be poisoned by impurities (As, S, Hg compounds). 3. Difficult to set up and maintain in routine laboratory work; calomel or Ag/AgCl electrodes are preferred as secondary references.

Common Mistakes to Avoid

1. Confusing order and molecularity. Order is always determined experimentally; molecularity applies only to elementary steps. A complex reaction has an overall order but no overall molecularity.

2. Using stoichiometric coefficients as reaction order. For the reaction aA + bB → products, the rate law is NOT necessarily Rate = k[A]^a[B]^b. The order must be determined experimentally. Only for elementary reactions do order and stoichiometry match.

3. Forgetting units of the rate constant. Each order has different units for k: zero order (mol L⁻¹ s⁻¹), first order (s⁻¹), second order (L mol⁻¹ s⁻¹). Always write units — marks are deducted for missing units.

4. Sign errors in the Nernst equation. The correct form is E = E⁰ - (0.0591/n) log Q. Students often write + instead of -. Remember: products in numerator, reactants in denominator of Q, and subtract the term.

5. Misapplying the van't Hoff factor. For dissociation, i > 1 (more particles than expected). For association, i < 1 (fewer particles). Use the correct formula: i = 1 + alpha(n-1) for dissociation; i = 1 - (1 - 1/n)alpha for association. Do not mix them up.

6. Confusing physisorption and chemisorption trends with temperature. Physisorption decreases with increasing temperature (exothermic, weak forces). Chemisorption first increases (because activation energy is needed) then decreases (because it is exothermic overall). Drawing a single monotonic trend for chemisorption is incorrect.

Exam Checklist

• [ ] I can write the Raoult's law equations for volatile and non-volatile solutes and solve numericals on relative lowering of vapour pressure, boiling point elevation, freezing point depression, and osmotic pressure.
• [ ] I can calculate the van't Hoff factor and degree of dissociation/association from colligative property data.
• [ ] I can apply the Nernst equation to calculate cell EMF at non-standard conditions and determine spontaneity from E\_cell and delta G.
• [ ] I can use Kohlrausch's law to calculate limiting molar conductivity and degree of dissociation of weak electrolytes.
• [ ] I can write and apply integrated rate laws for zero-order, first-order, and pseudo first-order reactions, including half-life calculations.
• [ ] I can use the Arrhenius equation (both forms) to calculate activation energy, rate constant at a different temperature, and interpret ln k vs 1/T plots.
• [ ] I can distinguish between physisorption and chemisorption and apply the Freundlich and Langmuir isotherms.
• [ ] I can classify colloids, explain Tyndall effect, Brownian motion, electrophoresis, and apply the Hardy-Schulze rule to determine coagulating power.