Class 11 Notes

Class 11 Chemistry Notes Chapter 5 (States of matter) – Chemistry Part-I Book

Philoid

Philoid

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Chemistry Part-I
Alright class, let's begin our detailed study of Chapter 5, 'States of Matter', from your NCERT Class 11 Chemistry textbook. This chapter is crucial as it lays the foundation for understanding the physical properties of substances, which often feature in various government examinations. We'll focus on the key concepts, laws, and formulas.

Chapter 5: States of Matter - Detailed Notes for Government Exam Preparation

1. Introduction

  • Matter exists primarily in three states: Solid, Liquid, and Gas. Plasma is sometimes considered the fourth state (at very high temperatures).
  • The state of a substance is determined by the balance between two opposing factors:
    • Intermolecular Forces (IMFs): Attractive forces between molecules (van der Waals forces, dipole-dipole interactions, hydrogen bonding). These tend to keep molecules together.
    • Thermal Energy: Energy possessed by molecules due to their motion (temperature is a measure of average kinetic energy). This tends to make molecules move apart.
  • Order of IMFs: Solid > Liquid > Gas
  • Order of Thermal Energy: Gas > Liquid > Solid

2. The Gaseous State

  • Characterized by negligible intermolecular forces (ideally) and high thermal energy.

  • Properties: No fixed shape or volume, highly compressible, exert pressure uniformly in all directions, mix evenly without mechanical aid (diffusion), low density.

  • Measurable Properties of Gases:

    • Mass (m): Usually expressed in grams (g) or kilograms (kg). Often represented by the number of moles (n). (n = mass / Molar Mass)
    • Volume (V): Space occupied by the gas. Units: Litre (L), millilitre (mL), cubic centimetre (cm³), cubic metre (m³). (1 L = 1 dm³ = 1000 mL = 1000 cm³; 1 m³ = 1000 L)
    • Pressure (P): Force exerted per unit area. Units: Pascal (Pa - SI unit), atmosphere (atm), bar, millimetres of mercury (mm Hg), torr. (1 atm = 101325 Pa = 1.01325 bar = 760 mm Hg = 760 torr; 1 bar = 10⁵ Pa) Measured using a manometer or barometer.
    • Temperature (T): Degree of hotness or coldness. Always use the absolute scale (Kelvin, K) in gas law calculations. K = °C + 273.15 (often approximated as 273).

  • Gas Laws: Describe the relationship between P, V, T, and n for gases.

    • a) Boyle's Law (Pressure-Volume Relationship): At constant temperature (T) and number of moles (n), the pressure (P) of a fixed amount of gas varies inversely with its volume (V).

      • Mathematically: P ∝ 1/V or PV = k (constant)
      • For two states: P₁V₁ = P₂V₂
      • Graphical Representation: P vs V graph is a rectangular hyperbola (isotherm). P vs 1/V graph is a straight line passing through the origin.

    • b) Charles's Law (Temperature-Volume Relationship): At constant pressure (P) and number of moles (n), the volume (V) of a fixed amount of gas is directly proportional to its absolute temperature (T).

      • Mathematically: V ∝ T or V/T = k' (constant)
      • For two states: V₁/T₁ = V₂/T₂
      • Absolute Zero: Temperature at which the volume of a gas theoretically becomes zero (-273.15 °C or 0 K).
      • Graphical Representation: V vs T (in K) graph is a straight line passing through the origin (isobar).

    • c) Gay-Lussac's Law (Pressure-Temperature Relationship): At constant volume (V) and number of moles (n), the pressure (P) of a fixed amount of gas is directly proportional to its absolute temperature (T).

      • Mathematically: P ∝ T or P/T = k'' (constant)
      • For two states: P₁/T₁ = P₂/T₂
      • Graphical Representation: P vs T (in K) graph is a straight line passing through the origin (isochore).

    • d) Avogadro's Law (Volume-Amount Relationship): At constant temperature (T) and pressure (P), the volume (V) of a gas is directly proportional to the number of moles (n).

      • Mathematically: V ∝ n or V/n = k''' (constant)
      • Conclusion: Equal volumes of all gases under the same conditions of temperature and pressure contain an equal number of molecules.
      • Molar Volume: Volume occupied by 1 mole of any gas at STP (Standard Temperature and Pressure: 273.15 K, 1 bar) is 22.71 L/mol. At older STP definition (273.15 K, 1 atm), it's 22.4 L/mol. Be mindful of the definition used in the exam.

  • Ideal Gas Equation: Combination of Boyle's, Charles's, and Avogadro's Laws.

    • PV = nRT
    • R = Universal Gas Constant. Its value depends on the units used for P, V, and T.
      • R = 8.314 J K⁻¹ mol⁻¹ (SI unit, when P in Pa, V in m³)
      • R = 0.0821 L atm K⁻¹ mol⁻¹ (when P in atm, V in L)
      • R ≈ 2 cal K⁻¹ mol⁻¹
    • An ideal gas is a hypothetical gas that perfectly obeys the ideal gas equation under all conditions.

  • Combined Gas Law: For a fixed amount of gas (n = constant).

    • (P₁V₁)/T₁ = (P₂V₂)/T₂

  • Density and Molar Mass Relationship:

    • From PV = nRT, substitute n = m/M (mass/Molar Mass)
    • PV = (m/M)RT => P = (m/V)(RT/M)
    • Since density (d) = m/V, P = d(RT/M) => d = PM/RT
    • Higher molar mass or higher pressure leads to higher density. Higher temperature leads to lower density.

  • Dalton's Law of Partial Pressures: The total pressure exerted by a mixture of non-reacting gases is equal to the sum of the partial pressures that each gas would exert if it were present alone in the same volume and at the same temperature.

    • P_total = p₁ + p₂ + p₃ + ... (at constant T, V)
    • Partial Pressure (pᵢ): Pressure exerted by an individual gas in the mixture.
    • pᵢ = xᵢ * P_total , where xᵢ is the mole fraction of gas 'i'.
    • xᵢ = (moles of gas i) / (total moles of all gases)
    • Application: Collection of gas over water. P_dry gas = P_total - Aqueous Tension (Aqueous tension is the partial pressure of water vapour at that temperature).

  • Graham's Law of Diffusion/Effusion: The rate of diffusion (mixing of gases) or effusion (escape of gas through a tiny hole) is inversely proportional to the square root of its molar mass (M) or density (d) at constant temperature and pressure.

    • Rate (r) ∝ 1/√M or Rate (r) ∝ 1/√d
    • For two gases A and B: r_A / r_B = √(M_B / M_A) = √(d_B / d_A)

  • Kinetic Molecular Theory of Gases (KMT): Explains the behaviour of ideal gases based on assumptions about molecular motion.

    • Postulates:
      1. Gases consist of a large number of tiny particles (atoms/molecules) that are far apart relative to their size.
      2. The actual volume of the gas molecules is negligible compared to the total volume of the container.
      3. Gas molecules are in constant, random motion, colliding with each other and the walls of the container.
      4. Collisions are perfectly elastic (no loss of kinetic energy).
      5. There are no significant attractive or repulsive forces between gas molecules (IMFs are negligible).
      6. The average kinetic energy of the gas molecules is directly proportional to the absolute temperature (T). KE_avg ∝ T.
    • Kinetic Gas Equation: P = (1/3)mnū²/V (where m = mass of one molecule, n = number of molecules, ū² = mean square speed, V = volume). Not usually needed for direct calculation in exams, but understanding is key.
    • Kinetic Energy: Average KE per molecule = (3/2)kT (k = Boltzmann constant = R/N_A). Average KE per mole = (3/2)RT. Importantly, at a given temperature, all gases have the same average kinetic energy.
    • Molecular Speeds: Molecules move at different speeds.
      • Root Mean Square Speed (u_rms): √(3RT/M)
      • Average Speed (u_av): √(8RT/πM)
      • Most Probable Speed (u_mp): √(2RT/M)
      • Relationship: u_rms > u_av > u_mp (Ratio ≈ 1.224 : 1.128 : 1)

  • Behaviour of Real Gases: Deviation from Ideal Behaviour

    • Real gases obey ideal gas laws only at LOW pressure and HIGH temperature.
    • Deviations occur at HIGH pressure and LOW temperature because:
      • Intermolecular forces become significant (Assumption 5 of KMT fails).
      • Molecular volume is no longer negligible compared to container volume (Assumption 2 of KMT fails).
    • Compressibility Factor (Z): Measures the deviation of a real gas from ideal behaviour.
      • Z = PV / nRT
      • For Ideal Gas: Z = 1 under all conditions.
      • For Real Gases:
        • Z < 1 (Negative deviation): Attractive forces dominate. Gas is more compressible than ideal. Occurs at low pressures (initially).
        • Z > 1 (Positive deviation): Repulsive forces dominate (due to molecular volume). Gas is less compressible than ideal. Occurs at high pressures.
    • Van der Waals Equation: Modification of the ideal gas equation for real gases.
      • [P + a(n/V)²] [V - nb] = nRT
      • For 1 mole: [P + a/V²] [V - b] = RT
      • 'a': Correction factor for intermolecular forces (Units: atm L² mol⁻² or Pa m⁶ mol⁻²). Higher 'a' means stronger IMFs.
      • 'b': Correction factor for molecular volume (Excluded volume or co-volume) (Units: L mol⁻¹ or m³ mol⁻¹). 'b' is related to the actual size of molecules (approximately 4 times the actual molecular volume).
    • Liquefaction of Gases: Gases can be liquefied by increasing pressure and decreasing temperature.
      • Critical Temperature (Tc): The maximum temperature above which a gas cannot be liquefied, no matter how high the pressure applied. Tc = 8a / 27Rb
      • Critical Pressure (Pc): The minimum pressure required to liquefy a gas at its critical temperature. Pc = a / 27b²
      • Critical Volume (Vc): The volume occupied by one mole of the gas at its critical temperature and pressure. Vc = 3b
      • Gases with higher Tc (due to stronger IMFs, i.e., higher 'a') are easier to liquefy.
    • Boyle Temperature (Tb): The temperature at which a real gas behaves like an ideal gas over an appreciable pressure range (Z ≈ 1). Tb = a / Rb.

3. The Liquid State

  • Intermediate state between gas and solid. Molecules are close together but can move past each other.

  • Properties: Definite volume, no definite shape, much less compressible than gases, higher density than gases, can diffuse (slower than gases).

  • Key Properties:

    • a) Vapour Pressure: The pressure exerted by the vapour molecules of a liquid in equilibrium with the liquid at a given temperature in a closed container.
      • Factors affecting: Nature of liquid (weaker IMFs -> higher vapour pressure), Temperature (higher T -> higher vapour pressure).
      • Independent of the amount of liquid or surface area (as long as some liquid is present).
    • b) Boiling Point: The temperature at which the vapour pressure of a liquid becomes equal to the external pressure.
      • Normal Boiling Point: Boiling point at 1 atm external pressure.
      • Standard Boiling Point: Boiling point at 1 bar external pressure.
      • Higher external pressure -> Higher boiling point. Lower external pressure (e.g., at high altitudes) -> Lower boiling point.
    • c) Surface Tension (γ): The force acting per unit length perpendicular to an imaginary line drawn on the surface of a liquid. It's the energy required to increase the surface area by one unit.
      • Cause: Unbalanced intermolecular forces experienced by molecules at the surface compared to those in the bulk. Liquids tend to minimize their surface area (e.g., spherical drops).
      • Units: N m⁻¹ or J m⁻².
      • Factors affecting: Nature of liquid (stronger IMFs -> higher surface tension), Temperature (higher T -> lower surface tension).
      • Consequences: Capillary rise or fall, spherical shape of drops/bubbles.
    • d) Viscosity (η): The measure of resistance to flow. "Thickness" of a fluid.
      • Cause: Intermolecular forces holding layers of liquid together.
      • Units: N s m⁻² or Pa s (SI unit). Common unit: poise (1 poise = 0.1 Pa s).
      • Factors affecting: Nature of liquid (stronger IMFs, complex shapes -> higher viscosity), Temperature (higher T -> lower viscosity for liquids; opposite for gases).

4. The Solid State (Briefly mentioned here, detailed in Class 12)

  • Particles are held in fixed positions, only vibrate about mean positions.
  • Definite shape and volume, high density, incompressible, rigid.

Multiple Choice Questions (MCQs)

  1. According to Boyle's Law, if the pressure of a fixed amount of gas is doubled at constant temperature, its volume becomes:
    a) Doubled
    b) Halved
    c) Remains constant
    d) Four times

  2. The value of the universal gas constant (R) depends on:
    a) Temperature of the gas
    b) Volume of the gas
    c) Number of moles of the gas
    d) Units of pressure, volume, and temperature

  3. At the same temperature and pressure, which of the following gases will have the highest rate of diffusion?
    a) O₂ (Molar mass = 32 g/mol)
    b) N₂ (Molar mass = 28 g/mol)
    c) CH₄ (Molar mass = 16 g/mol)
    d) CO₂ (Molar mass = 44 g/mol)

  4. The compressibility factor (Z) for an ideal gas is:
    a) Zero
    b) Less than 1
    c) Greater than 1
    d) Equal to 1

  5. The van der Waals equation constant 'a' is a measure of:
    a) Mean velocity of molecules
    b) Intermolecular forces of attraction
    c) Volume occupied by molecules
    d) Mean free path

  6. Critical temperature (Tc) is the temperature:
    a) At which a gas liquefies at 1 atm pressure
    b) Below which a gas can be liquefied by applying pressure
    c) Above which a gas cannot be liquefied, however high the pressure
    d) At which volume of gas becomes zero

  7. A gas collected over water has a total pressure of 750 mm Hg at 25°C. If the aqueous tension at 25°C is 24 mm Hg, what is the partial pressure of the dry gas?
    a) 774 mm Hg
    b) 750 mm Hg
    c) 726 mm Hg
    d) 24 mm Hg

  8. Which property of a liquid decreases with an increase in temperature?
    a) Vapour Pressure
    b) Boiling Point
    c) Surface Tension
    d) None of the above

  9. According to the Kinetic Theory of Gases, the average kinetic energy of gas molecules is directly proportional to:
    a) Pressure
    b) Volume
    c) Absolute Temperature
    d) Molar Mass

  10. Under which conditions do real gases deviate most significantly from ideal behaviour?
    a) High temperature and low pressure
    b) High temperature and high pressure
    c) Low temperature and low pressure
    d) Low temperature and high pressure

Answer Key for MCQs:

  1. b) Halved
  2. d) Units of pressure, volume, and temperature
  3. c) CH₄ (Lowest molar mass)
  4. d) Equal to 1
  5. b) Intermolecular forces of attraction
  6. c) Above which a gas cannot be liquefied, however high the pressure
  7. c) 726 mm Hg (P_dry = P_total - Aqueous Tension = 750 - 24)
  8. c) Surface Tension (Viscosity also decreases)
  9. c) Absolute Temperature
  10. d) Low temperature and high pressure

Remember to thoroughly understand the concepts behind these laws and definitions. Practice numerical problems based on the gas laws and the ideal gas equation. Good luck with your preparation!

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