A Level H1 Physics Thermal Physics Quiz

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A-Level Physics H1 Quiz - Thermal Physics

Name: ____________________ Class: ____________________ Date: ____________________ Score: ________ / 50

Duration: 60 Minutes Total Marks: 50 Marks

Instructions:

1. Answer all questions.
2. Show all necessary working for calculation questions.
3. Use $g = 9.81 \text{ m s}^{-2}$ and the standard values for constants unless otherwise stated.
4. Write your answers in the spaces provided.

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Section A: Fundamental Concepts (Questions 1-5)

1. Define the term specific heat capacity. [2]
   
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2. State the difference between internal energy and thermal energy. [2]
   
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3. A metal block of mass $0.50 \text{ kg}$ and specific heat capacity $385 \text{ J kg}^{-1} \text{ K}^{-1}$ is heated from $20^\circ\text{C}$ to $80^\circ\text{C}$. Calculate the energy supplied. [2]
   
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4. Explain why the temperature of a substance remains constant during a change of state (e.g., melting). [2]
   
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5. State the First Law of Thermodynamics in terms of energy conservation. [2]
   
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Section B: Calculations and Applications (Questions 6-15)

6. An electric heater of power $100 \text{ W}$ is used to heat $200 \text{ g}$ of water. Calculate the time taken to raise the temperature by $10^\circ\text{C}$. (Specific heat capacity of water = $4180 \text{ J kg}^{-1} \text{ K}^{-1}$) [3]
   
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7. A $0.10 \text{ kg}$ piece of copper at $100^\circ\text{C}$ is dropped into $0.20 \text{ kg}$ of water at $20^\circ\text{C}$. Calculate the final equilibrium temperature. (Copper $c = 390 \text{ J kg}^{-1} \text{ K}^{-1}$, Water $c = 4180 \text{ J kg}^{-1} \text{ K}^{-1}$) [3]
   
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8. Define specific latent heat of fusion and state its SI unit. [2]
   
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9. Calculate the energy required to melt $50 \text{ g}$ of ice at $0^\circ\text{C}$. (Specific latent heat of fusion of ice = $3.34 \times 10^5 \text{ J kg}^{-1}$) [2]
   
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10. A sample of gas is compressed isothermally. Describe what happens to the internal energy and the work done on the gas. [3]
    
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11. An ideal gas undergoes an adiabatic expansion. Explain why the temperature of the gas decreases. [3]
    
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12. A $2.0 \text{ mol}$ sample of an ideal gas is kept at $300 \text{ K}$. Calculate the total internal energy of the gas. (Assume monatomic gas, $R = 8.31 \text{ J mol}^{-1} \text{ K}^{-1}$) [3]
    
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13. A cylinder contains an ideal gas at pressure $1.0 \times 10^5 \text{ Pa}$ and volume $2.0 \times 10^{-3} \text{ m}^3$. If the volume is halved at constant temperature, determine the new pressure. [2]
    
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14. Explain the relationship between the average kinetic energy of molecules in an ideal gas and the absolute temperature. [2]
    
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15. A $0.5 \text{ kg}$ block of aluminum is heated. If the energy supplied is $2000 \text{ J}$ and the temperature rises by $15 \text{ K}$, calculate the specific heat capacity of aluminum. [2]
    
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Section C: Advanced Analysis (Questions 16-20)

16. A gas expands from volume $V_1$ to $V_2$ against a constant external pressure $P$. Derive an expression for the work done by the gas. [3]
    
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17. Compare the rate of heat transfer by conduction in a metal rod versus a plastic rod of the same dimensions. Explain your answer using the concept of free electrons. [3]
    
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18. A $100 \text{ g}$ piece of ice at $-10^\circ\text{C}$ is heated until it becomes water at $20^\circ\text{C}$. Calculate the total heat energy required. (Ice $c = 2100 \text{ J kg}^{-1} \text{ K}^{-1}$, $L_f = 3.34 \times 10^5 \text{ J kg}^{-1}$, Water $c = 4180 \text{ J kg}^{-1} \text{ K}^{-1}$) [4]
    
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19. Describe the process of evaporation and explain why it leads to the cooling of the remaining liquid. [3]
    
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20. A gas is heated at constant pressure. Explain why the work done by the gas is equal to the change in internal energy minus the heat added (or use the 1st Law to explain the energy split). [3]
    
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Answer Key - A-Level Physics H1 Quiz: Thermal Physics

1. Definition: The amount of heat energy required to raise the temperature of unit mass of a substance by one Kelvin (or one degree Celsius). [2]

2. Difference: Internal energy is the sum of the random kinetic and potential energies of the molecules. Thermal energy is the energy transferred between systems due to a temperature difference. [2]

3. $Q = mc\Delta T = 0.50 \times 385 \times (80 - 20) = 0.50 \times 385 \times 60 = 11,550 \text{ J}$. [2]

4. Explanation: During a change of state, the energy supplied is used to break the intermolecular bonds (increasing potential energy) rather than increasing the average kinetic energy of the molecules. [2]

5. 1st Law: $\Delta U = Q - W$ (or $Q = \Delta U + W$). The change in internal energy of a system is equal to the heat added to the system minus the work done by the system on its surroundings. [2]

6. $Q = mc\Delta T = 0.20 \times 4180 \times 10 = 8360 \text{ J}$. $t = Q / P = 8360 / 100 = 83.6 \text{ s}$. [3]

7. Heat lost by copper = Heat gained by water. $m_c c_c (100 - T) = m_w c_w (T - 20)$ $0.10 \times 390 \times (100 - T) = 0.20 \times 4180 \times (T - 20)$ $39(100 - T) = 836(T - 20)$ $3900 - 39T = 836T - 16720$ $875T = 20620 \Rightarrow T \approx 23.6^\circ\text{C}$. [3]

8. Definition: The energy required to change unit mass of a substance from solid to liquid at a constant temperature. Unit: $\text{J kg}^{-1}$. [2]

9. $Q = mL_f = 0.050 \times 3.34 \times 10^5 = 16,700 \text{ J}$. [2]

10. Isothermal: $\Delta T = 0 \Rightarrow \Delta U = 0$. Since $\Delta U = Q - W$, then $Q = W$. The heat added to the gas is exactly equal to the work done by the gas. [3]

11. Adiabatic: $Q = 0$. $\Delta U = -W$. As the gas expands, it does work on the surroundings. This energy comes from the internal energy of the gas, causing $\Delta U$ to decrease and thus temperature to drop. [3]

12. For monatomic ideal gas: $U = \frac{3}{2} nRT$. $U = 1.5 \times 2.0 \times 8.31 \times 300 = 7479 \text{ J}$. [3]

13. $P_1 V_1 = P_2 V_2$ (Boyle's Law). $1.0 \times 10^5 \times 2.0 \times 10^{-3} = P_2 \times (1.0 \times 10^{-3})$ $P_2 = 2.0 \times 10^5 \text{ Pa}$. [2]

14. The absolute temperature is directly proportional to the average kinetic energy of the molecules ($KE_{avg} = \frac{3}{2} k_B T$). [2]

15. $c = Q / (m\Delta T) = 2000 / (0.5 \times 15) = 2000 / 7.5 = 266.7 \text{ J kg}^{-1} \text{ K}^{-1}$. [2]

16. $W = \int P dV$. For constant pressure: $W = P(V_2 - V_1)$. [3]

17. Comparison: Metal rod has a much higher rate of heat transfer. Explanation: Metals possess free electrons that can move rapidly through the lattice, transferring kinetic energy much faster than the vibrational modes (phonons) alone in plastic. [3]

18. $Q_1$ (warm ice): $0.1 \times 2100 \times 10 = 2100 \text{ J}$ $Q_2$ (melt): $0.1 \times 3.34 \times 10^5 = 33,400 \text{ J}$ $Q_3$ (warm water): $0.1 \times 4180 \times 20 = 8360 \text{ J}$ Total $Q = 2100 + 33400 + 8360 = 43,860 \text{ J}$. [4]

19. Process: Molecules with higher than average kinetic energy escape from the surface of the liquid. Cooling: Since the highest energy molecules leave, the average kinetic energy of the remaining molecules decreases, leading to a lower temperature. [3]

20. From 1st Law: $Q = \Delta U + W$. When heated at constant pressure, the energy supplied ($Q$) is split: some increases the internal energy (raising temperature) and some is used to do work as the gas expands. [3]