Class 12 Chemistry Notes Chapter 3 (Thermochemical Measurement) – Lab Manual (English) Book

Detailed Notes with MCQs of Chapter 3: Thermochemical Measurement from your Chemistry Lab Manual. This is a crucial topic, not just for your practical exams but also as it forms the basis for understanding energy changes in chemical reactions, which frequently appears in various government entrance exams. Pay close attention to the principles, procedures, and calculations involved.

Thermochemical Measurement: Key Concepts & Principles

1. Thermochemistry: The branch of chemistry concerned with the heat changes (energy in the form of heat) that accompany chemical reactions and physical transformations.

2. System & Surroundings:

   • System: The specific part of the universe being studied (e.g., the reactants and products in a beaker).
   • Surroundings: Everything else in the universe outside the system (e.g., the beaker, the air, the lab bench).
   • In calorimetry experiments, the system is usually the chemical reaction (dissolution, neutralization), and the surroundings are the solvent (water), the calorimeter, and the thermometer.

3. Enthalpy (H): A thermodynamic property representing the total heat content of a system at constant pressure. We cannot measure absolute enthalpy, but we can measure the change in enthalpy (ΔH).

4. Change in Enthalpy (ΔH): The heat absorbed or released by a system during a process occurring at constant pressure.

   • ΔH = H_products - H_reactants
   • Exothermic Reaction: Releases heat to the surroundings. The system loses heat, so ΔH is negative (ΔH < 0). The temperature of the surroundings (e.g., water in the calorimeter) increases.
   • Endothermic Reaction: Absorbs heat from the surroundings. The system gains heat, so ΔH is positive (ΔH > 0). The temperature of the surroundings decreases.

5. Calorimetry: The science of measuring heat changes.

   • Calorimeter: An insulated device used to measure the heat absorbed or released during a chemical or physical process. A simple lab calorimeter often consists of two nested polystyrene (Styrofoam) cups with a lid, a stirrer, and a thermometer. Polystyrene is used for its good insulating properties.
   • Principle: Based on the law of conservation of energy. Heat released by the reaction = Heat absorbed by the surroundings (solution + calorimeter), or Heat absorbed by the reaction = Heat lost by the surroundings.
     • q_reaction = - (q_solution + q_calorimeter)

6. Heat Capacity (C) & Specific Heat Capacity (c):

   • Heat Capacity (C): The amount of heat required to raise the temperature of a substance (or a device like a calorimeter) by 1°C or 1 K. Unit: J/°C or J/K.
   • Specific Heat Capacity (c): The amount of heat required to raise the temperature of 1 gram of a substance by 1°C or 1 K. Unit: J/(g·°C) or J/(g·K). For water, c ≈ 4.18 J/(g·°C).

7. Calculating Heat Change (q):

   • q = m × c × ΔT (where m = mass of substance, c = specific heat capacity, ΔT = change in temperature = T_final - T_initial)
   • q = C × ΔT (where C = heat capacity of the device/system)

Common Thermochemical Measurements in the Lab

A. Enthalpy of Dissolution (Δ_sol_H)

• Definition: The enthalpy change when one mole of a substance dissolves in a specified amount of solvent (usually a large excess, so it approaches infinite dilution) at constant temperature and pressure.
• Procedure:
  1. Measure a known volume (hence mass, since density of water ≈ 1 g/mL) of water into the calorimeter.
  2. Record the initial steady temperature of the water (T_initial).
  3. Weigh accurately a known mass of the solute (e.g., CuSO₄·5H₂O, KCl, NaOH).
  4. Add the solute to the water in the calorimeter, stir gently but continuously to dissolve it completely and distribute the heat evenly.
  5. Record the maximum (for exothermic) or minimum (for endothermic) temperature reached (T_final).
• Calculations:
  1. Calculate the change in temperature: ΔT = T_final - T_initial.
  2. Calculate the total mass of the solution: m_solution = mass of water + mass of solute.
  3. Calculate the heat absorbed or released by the solution: q_solution = m_solution × c_solution × ΔT. (Assume c_solution ≈ c_water = 4.18 J/(g·°C)).
  4. Calculate the heat absorbed or released by the calorimeter: q_calorimeter = C_calorimeter × ΔT. (Often, the heat capacity of a simple polystyrene calorimeter is small and may be neglected, or it can be determined experimentally).
  5. Calculate the total heat change for the reaction: q_reaction = - (q_solution + q_calorimeter). (If neglecting q_calorimeter, q_reaction = -q_solution).
  6. Calculate the moles of solute dissolved: moles = mass of solute / molar mass of solute.
  7. Calculate the enthalpy of dissolution per mole: Δ_sol_H = q_reaction / moles of solute. (Ensure units are in kJ/mol).
• Sign: If T_final > T_initial (exothermic), ΔT is positive, q_solution is positive, q_reaction is negative, Δ_sol_H is negative. If T_final < T_initial (endothermic), ΔT is negative, q_solution is negative, q_reaction is positive, Δ_sol_H is positive.
• Examples: Dissolution of NaOH, H₂SO₄, anhydrous CaCl₂, CuSO₄·5H₂O is typically exothermic. Dissolution of KNO₃, NH₄Cl, KCl is typically endothermic.

B. Enthalpy of Neutralization (Δ_neut_H)

• Definition: The enthalpy change when one mole of water is formed from the reaction between an acid and a base under standard conditions (usually 1 M solutions, 298 K, 1 atm).
  • H⁺(aq) + OH⁻(aq) → H₂O(l)
• Strong Acid + Strong Base: The reaction is essentially the formation of water from H⁺ and OH⁻ ions. The enthalpy change is nearly constant, approximately -57.1 kJ/mol. Examples: HCl + NaOH, HNO₃ + KOH.
• Weak Acid/Base Involvement: If either the acid or the base (or both) is weak, some energy is used to ionize the weak electrolyte. Therefore, the heat released is less than 57.1 kJ/mol (i.e., Δ_neut_H is less negative). Example: CH₃COOH + NaOH.
• Procedure:
  1. Measure equal volumes of acid and base solutions of known concentration (e.g., 50 mL of 1 M HCl and 50 mL of 1 M NaOH).
  2. Record their initial temperatures separately. Ideally, they should be the same; if not, take the average as T_initial.
  3. Pour one solution into the calorimeter, then quickly add the other. Stir gently.
  4. Record the maximum temperature reached (T_final).
• Calculations:
  1. Calculate the change in temperature: ΔT = T_final - T_initial.
  2. Calculate the total mass of the mixture: m_mixture = volume of acid + volume of base (assuming density ≈ 1 g/mL).
  3. Calculate the heat released: q_solution = m_mixture × c_solution × ΔT. (Assume c_solution ≈ c_water = 4.18 J/(g·°C)).
  4. Account for q_calorimeter if necessary: q_reaction = - (q_solution + q_calorimeter).
  5. Calculate the moles of acid and base reacted. Determine the moles of water formed (usually equal to moles of acid/base if stoichiometry is 1:1 and concentrations/volumes are equal, or based on the limiting reactant).
  6. Calculate the enthalpy of neutralization per mole of water formed: Δ_neut_H = q_reaction / moles of H₂O formed. (Ensure units are in kJ/mol).

C. Hess's Law of Constant Heat Summation

• Statement: The total enthalpy change for a chemical reaction is independent of the pathway taken, provided the initial and final conditions are the same. It depends only on the initial and final states.
• Application: Allows calculation of enthalpy changes for reactions that are difficult or impossible to measure directly (e.g., very slow reactions, reactions with side products). This is done by manipulating (reversing, multiplying) known enthalpy changes of related reactions and adding them algebraically.
• Example: Calculating ΔH for C(s) + ½ O₂(g) → CO(g).
  We know:
  (1) C(s) + O₂(g) → CO₂(g) ΔH₁ = -393.5 kJ
  (2) CO(g) + ½ O₂(g) → CO₂(g) ΔH₂ = -283.0 kJ
  Reverse reaction (2) and change the sign of ΔH₂:
  (2') CO₂(g) → CO(g) + ½ O₂(g) ΔH₂' = +283.0 kJ
  Add reaction (1) and (2'):
  C(s) + O₂(g) + CO₂(g) → CO₂(g) + CO(g) + ½ O₂(g)
  Simplifying gives: C(s) + ½ O₂(g) → CO(g)
  The enthalpy change is: ΔH = ΔH₁ + ΔH₂' = (-393.5) + (+283.0) = -110.5 kJ/mol.

Important Precautions for Calorimetry Experiments:

1. Use a well-insulated calorimeter (like polystyrene cups with a lid) to minimize heat exchange with the surroundings.
2. Measure volumes and masses accurately.
3. Use a sensitive thermometer (preferably calibrated to 0.1 °C).
4. Stir the contents continuously but gently for uniform temperature distribution.
5. Record the initial temperature only after it becomes constant.
6. Record the final temperature accurately (maximum for exothermic, minimum for endothermic). Extrapolating the cooling/warming curve can improve accuracy if the temperature change is slow.
7. Add the solute or reactant quickly to minimize heat loss during addition.
8. Ensure complete dissolution or reaction.
9. Handle chemicals like concentrated acids, bases (NaOH, KOH), and hygroscopic substances with appropriate safety measures (gloves, goggles).

Summary of Calculations:

• Heat Change (q): q = m × c × ΔT (for solution) or q = C × ΔT (for calorimeter)
• Reaction Heat: q_reaction = - (q_solution + q_calorimeter)
• Molar Enthalpy Change (ΔH): ΔH = q_reaction / moles (of solute dissolved or water formed)
• Units: Pay close attention to units (Joules vs. Kilojoules, grams, moles, °C or K for ΔT). Final ΔH is usually expressed in kJ/mol.
• Sign Convention: Negative ΔH for exothermic, Positive ΔH for endothermic.

This covers the essential aspects of thermochemical measurements relevant to your lab manual and competitive exams. Ensure you understand the underlying principles and practice the calculations.

Multiple Choice Questions (MCQs)

1. The enthalpy change accompanying a reaction at constant pressure is denoted by:
   a) ΔU
   b) ΔH
   c) q
   d) w

2. In an endothermic process:
   a) Heat is released, ΔH is positive
   b) Heat is absorbed, ΔH is positive
   c) Heat is released, ΔH is negative
   d) Heat is absorbed, ΔH is negative

3. A simple calorimeter, like nested polystyrene cups, works on the principle of:
   a) Minimizing pressure changes
   b) Maximizing heat exchange with surroundings
   c) Minimizing heat exchange with surroundings
   d) Measuring volume changes accurately

4. When 5.85 g of NaCl (Molar mass = 58.5 g/mol) is dissolved in 100 mL of water, the temperature drops by 1.2 °C. Assuming the specific heat capacity of the solution is 4.18 J/(g·°C) and density is 1 g/mL, what is the approximate enthalpy of solution (Δ_sol_H) of NaCl? (Neglect calorimeter heat capacity)
   a) +5.3 kJ/mol
   b) -5.3 kJ/mol
   c) +530 J/mol
   d) -530 J/mol

5. The enthalpy of neutralization of a strong acid (like HCl) with a strong base (like NaOH) is approximately -57.1 kJ/mol. This value corresponds to the reaction:
   a) HCl(aq) + NaOH(aq) → NaCl(aq) + H₂O(l)
   b) H⁺(aq) + OH⁻(aq) → H₂O(l)
   c) H⁺(aq) + Cl⁻(aq) → HCl(aq)
   d) Na⁺(aq) + OH⁻(aq) → NaOH(aq)

6. If the enthalpy of neutralization of HCN (a weak acid) with NaOH (a strong base) is -12.1 kJ/mol, what is the enthalpy of ionization of HCN? (Use Δ_neut_H for strong acid/base = -57.1 kJ/mol)
   a) -45.0 kJ/mol
   b) +45.0 kJ/mol
   c) -69.2 kJ/mol
   d) +69.2 kJ/mol

7. Which law states that the total enthalpy change for a reaction is independent of the pathway taken?
   a) First Law of Thermodynamics
   b) Second Law of Thermodynamics
   c) Hess's Law
   d) Le Chatelier's Principle

8. During the determination of enthalpy of dissolution, stirring is essential to:
   a) Increase the rate of dissolution only
   b) Ensure uniform temperature distribution
   c) Prevent heat loss
   d) Measure the initial temperature accurately

9. Which of the following dissolutions is typically endothermic?
   a) Concentrated H₂SO₄ in water
   b) Anhydrous CaCl₂ in water
   c) NaOH pellets in water
   d) NH₄Cl crystals in water

10. In the calculation q = m × c × ΔT, 'm' represents:
    a) Mass of solute only
    b) Mass of solvent only
    c) Mass of solution (solute + solvent)
    d) Molar mass of solute

Answer Key for MCQs:

1. b
2. b
3. c
4. a (Calculation: moles NaCl = 5.85/58.5 = 0.1 mol. m_solution ≈ 100g + 5.85g ≈ 105.85g. q_solution = 105.85g * 4.18 J/(g·°C) * (-1.2 °C) ≈ -531 J. q_reaction = -q_solution ≈ +531 J. Δ_sol_H = +531 J / 0.1 mol = +5310 J/mol = +5.31 kJ/mol)
5. b
6. b (Explanation: ΔH(ionization) + ΔH(neutralization of H⁺+OH⁻) = ΔH(neutralization weak acid). ΔH(ionization) + (-57.1 kJ/mol) = -12.1 kJ/mol. ΔH(ionization) = -12.1 + 57.1 = +45.0 kJ/mol)
7. c
8. b
9. d
10. c