Thermodynamics Practice Questions I with Answers

1. A chemical reaction that absorbs heat from the surroundings is said to be __________ and has a __________ ∆H at constant pressure.
• A) endothermic, positive
• B) endothermic, negative
• C) exothermic, negative
• D) exothermic, positive
• E) exothermic, neutral
1. The value of ∆H for the reaction below is -126 kJ. How much heat (in kJ) is released when 2.00 mol of NaOH is formed in the reaction?

2Na2O2 (s) + 2H2O (l) ¬ 4NaOH (s) + O2 (g)

• A) 252
• B) -126
• C) 7.8
• D) 63
• E) 3.9
1. The specific heat of liquid bromine is 0.226 J/g-K. How much heat (J) is required to raise the temperature of 10.0 mL of bromine from 25.00ºC to 27.30ºC? The density of liquid bromine: 3.12 g/mL.
• A) 16.2
• B) 10.4
• C) 32.4
• D) 5.20
• E) 300
1. Which is true if ∆H is -95J?
• A) Both the system and the surroundings are gaining 95 J
• B) Both the system and the surroundings are losing 95 J
• C) The system is gaining 95J and the surroundings are losing 95 J
• D) The system is losing 95J and the surroundings are gaining 95 J
• E) The system is gaining 95J and the surroundings are gaining 95 J
1. Which statement about enthalpy is true?
• A) Heat is given off to the surroundings in endothermic reactions.
• B) Some substances have a negative specific heat capacity.
• C) Specific heat capacity is the same for all liquids.
• D) The sign of ∆H is always negative in exothermic reactions.
1. What happens to the value of ∆H for a thermochemical reaction if the reaction is reversed?
• A) ∆H has the same numerical value, and the sign changes.
• B) ∆H has the same numerical value, and the sign remains the same.
• C) ∆H is the reciprocal of the original value, and the sign changes.
• D) ∆H is the reciprocal of the original value, and the sign remains the same.
1. Which of the following processes is exothermic?
• A) Ether evaporating
• B) Ice melting
• C) Steam condensing
• D) Water decomposing
1. Calculate the work for the expansion of CO2 from 1.0 to 4.7 liters against a pressure of 1.0 atm at constant temperature.
• A) 3.7 L atm
• B) 4.7 L atm
• C) 0 L atm
• D) -3.7 L atm
• E) -4.7 L atm
1. Using the data below, what is the order of changes that occur when ethanol is heated from -25.0ºC to 85.0ºC?

Boiling point of ethanol: 78.5ºC                 Melting point of ethanol: -117.3ºC

• A) Phase change, temperature change
• B) Phase change, temperature change, phase change
• C) Temperature change, phase change
• D) Temperature change, phase change, temperature change

1. The value of ∆H for the reaction below is -336 kJ. Calculate the heat (kJ) released to the surroundings when 23.0 g of HCl is formed.

CH4 (g) + 3Cl2 (g) ¬ CHCl3 (l) + 3HCl (g)

• A) 211
• B) 177
• C) 70.7
• D) -336
• E) 2.57 x 103
1. Which of the following properties is (are) intensive properties?
• mass
• temperature
• . volume
• concentration
• energy
• A) I, III, and V
• B) II only
• C) II and IV
• D) III and IV
• E) I and V
1. How much heat is required to raise the temperature of a 5.75-g sample of iron (specific heat = 0.450 J/g°C) from 25.0°C to 79.8°C?
• A) 2.54 J
• B) 315 J
• C) 700 J
• D) 848 J
• E) 142 J
1. Two metals of equal mass with different heat capacities are subjected to the same amount of heat. Which undergoes the smallest change in temperature?
• A) The metal with the higher heat capacity.
• B) The metal with the lower heat capacity.
• C) Both undergo the same change in temperature.
• D) You need to know the initial temperatures of the metals.
• E) You need to know which metals you have.
1. A 32.5 g piece of aluminum (which has a molar heat capacity of 24.03 J/°C·mol) is heated to 82.4°C and dropped into a calorimeter containing water (specific heat capacity of water is 4.18 J/g°C) initially at 22.3°C. The final temperature of the water is 24.2°C. Ignoring significant figures, calculate the mass of water in the calorimeter.
• A) 212 g
• B) 5.72 kg
• C) 6.42 g
• D) 1.68 kg
• E) none of these
1. A 45.9 g sample of a metal is heated to 95.2°C and then placed in a calorimeter containing 120.0 g of water (c = 4.18 J/g°C) at 21.6°C. The final temperature of the water is 24.5°C. Which metal was used?
• A) Aluminum (c = 0.89 J/g°C)
• B) Iron (c = 0.45 J/g°C)
• C) Copper (c = 0.20 J/g°C)
• D) Lead (c = 0.14 J/g°C)
• E) none of these
1. You take 295.5 g of a solid at 30.0°C and let it melt in 425 g of water. The water temperature decreases from 85.1°C to 30.0°C. Calculate the heat of fusion of this solid.
• A) 160 J/g
• B) 166 J/g
• C) 331 J/g
• D) 721 J/g
• E) cannot solve without the heat capacity of the solid
1. The enthalpy of fusion of ice is 6.020 kJ/mol. The heat capacity of liquid water is 75.4 J/mol·°C. What is the smallest number of ice cubes at 0°C, each containing one mole of water, necessary to cool 500 g of liquid water initially at 20°C to 0°C?
• A) 1
• B) 7
• C) 14
• D) 15
• E) 126
1. 0 mL of pure water at 282 K is mixed with 50.0 mL of pure water at 306 K. What is the final temperature of the mixture?
• A) 294 K
• B) 297 K
• C) 342 K
• D) 588 K
• E) 24 K
1. Consider the reaction

H2(g) + O2(g) àH2O(l)     ∆H° = –286 kJ

Which of the following is true?

• A) The reaction is exothermic.
• B) The reaction is endothermic.
• C) The enthalpy of the products is less than that of the reactants.
• D) Heat is absorbed by the system.
• E) Both A and C are true.
1. What is the specific heat capacity of a metal if it requires 178.1 J to change the temperature of 15.0 g of the metal from 25.00°C to 32.00°C?
• A) 0.590 J/g°C
• B) 11.9 J/g°C
• C) 25.4 J/g°C
• D) 1.70 J/g°C
• E) 283 J/g°C

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