Math Problem Statement

1 mole of an ideal gas expands isothermally and reversibly from 5 dm^3 to 10 dm^3 at 300K. Calculate q, w, ∆U, ∆H, ∆G and ∆A.

Solution

To solve this problem, we will use the properties of an isothermal, reversible expansion of an ideal gas.

Given Data

  • Initial volume, V1=5dm3V_1 = 5 \, \text{dm}^3
  • Final volume, V2=10dm3V_2 = 10 \, \text{dm}^3
  • Temperature, T=300KT = 300 \, \text{K}
  • Number of moles, n=1moln = 1 \, \text{mol}
  • For an ideal gas, R=8.314J K1mol1R = 8.314 \, \text{J K}^{-1} \text{mol}^{-1}

Key Formulas and Concepts

  1. Isothermal Process: Since the temperature is constant, the internal energy change ΔU=0\Delta U = 0 and the enthalpy change ΔH=0\Delta H = 0 for an ideal gas.

  2. Heat and Work:

    • Heat absorbed by the system in an isothermal, reversible expansion q=nRTlnV2V1q = nRT \ln \frac{V_2}{V_1}.
    • The work done by the gas w=qw = -q in an isothermal process, as ΔU=0\Delta U = 0.
  3. Gibbs Free Energy Change (∆G):

    • For an isothermal, reversible process, ΔG=nRTlnV2V1\Delta G = nRT \ln \frac{V_2}{V_1}.
  4. Helmholtz Free Energy Change (∆A):

    • For an isothermal process, ΔA=w=nRTlnV2V1\Delta A = w = -nRT \ln \frac{V_2}{V_1}.

Solution Steps

1. Calculate qq (Heat Absorbed)

q=nRTlnV2V1q = nRT \ln \frac{V_2}{V_1} Substitute the values: q=(1mol)(8.314J K1mol1)(300K)ln105q = (1 \, \text{mol})(8.314 \, \text{J K}^{-1} \text{mol}^{-1})(300 \, \text{K}) \ln \frac{10}{5} q=2494.2J×ln2q = 2494.2 \, \text{J} \times \ln 2 q2494.2J×0.693q \approx 2494.2 \, \text{J} \times 0.693 q1729.6Jq \approx 1729.6 \, \text{J}

2. Calculate ww (Work Done)

Since w=qw = -q for an isothermal process: w=1729.6Jw = -1729.6 \, \text{J}

3. Calculate ΔU\Delta U (Change in Internal Energy)

For an ideal gas undergoing an isothermal process: ΔU=0\Delta U = 0

4. Calculate ΔH\Delta H (Change in Enthalpy)

For an ideal gas in an isothermal process: ΔH=0\Delta H = 0

5. Calculate ΔG\Delta G (Change in Gibbs Free Energy)

ΔG=q=1729.6J\Delta G = -q = -1729.6 \, \text{J}

6. Calculate ΔA\Delta A (Change in Helmholtz Free Energy)

ΔA=w=1729.6J\Delta A = w = -1729.6 \, \text{J}

Summary of Results

  • q=1729.6Jq = 1729.6 \, \text{J}
  • w=1729.6Jw = -1729.6 \, \text{J}
  • ΔU=0\Delta U = 0
  • ΔH=0\Delta H = 0
  • ΔG=1729.6J\Delta G = -1729.6 \, \text{J}
  • ΔA=1729.6J\Delta A = -1729.6 \, \text{J}

Would you like a deeper explanation of any specific step?


Here are some related questions to explore further:

  1. How would the calculations change if the expansion were adiabatic instead of isothermal?
  2. What would happen to ΔG\Delta G and ΔA\Delta A if the process were not reversible?
  3. Can you calculate the entropy change for this process?
  4. How does the ideal gas assumption affect these calculations?
  5. How would these values change if the temperature were different?

Tip: In isothermal processes for ideal gases, ΔU\Delta U and ΔH\Delta H are always zero since they depend only on temperature changes.

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Math Problem Analysis

Mathematical Concepts

Thermodynamics
Ideal Gas Laws
Isothermal Processes

Formulas

q = nRT ln(V2/V1)
w = -q
∆U = 0 (for isothermal process of an ideal gas)
∆H = 0 (for isothermal process of an ideal gas)
∆G = nRT ln(V2/V1)
∆A = -nRT ln(V2/V1)

Theorems

First Law of Thermodynamics
Properties of Isothermal Expansion in Ideal Gases

Suitable Grade Level

Undergraduate