## Thermodynamics Practice Questions II with Answers

1. If 5.0 kJ of energy is added to a 15.5-g sample of water at 10.°C, the water is
• A) boiling
• B) completely vaporized
• C) frozen solid
• D) decomposed
• E) still a liquid
1. A chunk of lead at 91.6°C was added to 200.0 g of water at 15.5°C. The specific heat of lead is 0.129 J/g°C, and the specific heat of water is 4.18 J/g°C. When the temperature stabilized, the temperature of the mixture was 17.9°C. Assuming no heat was lost to the surroundings, what was the mass of lead added?
• A) 1.57 kg
• B) 170 g
• C) 204 g
• D) 211 g
• E) none of these
1. On a cold winter day, a steel metal fence post feels colder than a wooden fence post of identical size because:
• A) The specific heat capacity of steel is higher than the specific heat capacity of wood.
• B) The specific heat capacity of steel is lower than the specific heat capacity of wood.
• C) Steel has the ability to resist a temperature change better than wood.
• D) The mass of steel is less than wood so it loses heat faster.
• E) Two of the above statements are true.
1. If a student performs an endothermic reaction in a calorimeter, how does the calculated value of ∆H differ from the actual value if the heat exchanged with the calorimeter is not taken into account?
• A) ∆Hcalc would be more negative because the calorimeter always absorbs heat from the reaction.
• B) ∆Hcalc would be less negative because the calorimeter would absorb heat from the reaction.
• C) ∆Hcalc would be more positive because the reaction absorbs heat from the calorimeter.
• D) ∆Hcalc would be less positive because the reaction absorbs heat from the calorimeter.
• E) ∆Hcalc would equal the actual value because the calorimeter does not absorb heat.
1. Consider the reaction:

When a 21.1-g sample of ethyl alcohol (molar mass = 46.07 g/mol) is burned, how much energy is released as heat?

• A) 0.458 kJ
• B) 0.627 kJ
• C) 6.27 x 102 kJ
• D) 2.89 x 104 kJ
• E) 2.18 kJ
1. CH4(g) + 4Cl2(g) à CCl4(g) + 4HCl(g), ∆H = –434 kJ

Based on the above reaction, what energy change occurs when 1.2 moles of methane (CH4) reacts?

• A) 5.2 x 105 J are released.
• B) 5.2 x 105 J are absorbed.
• C) 3.6 x 105 J are released.
• D) 3.6 x 105 J are absorbed.
• E) 4.4 x 105 J are released.
1. Given the equation S(s) + O2(g) à SO2(g), ∆H = –296 kJ, which of the following statement(s) is (are) true?
• The reaction is exothermic.
• When 0.500 mole sulfur is reacted, 148 kJ of energy is released.
• When 32.0 g of sulfur are burned, 2.96 x105 J of energy is released.
• A) All are true.
• B) None is true.
• C) I and II are true.
• D) I and III are true.
• E) Only II is true.
1. When 0.236 mol of a weak base (A) is reacted with excess HCl, 6.91 kJ of energy is released as heat. What is H for this reaction per mole of A consumed?
• A) –34.2 kJ
• B) –59.4 kJ
• C) –29.3 kJ
• D) 34.2 kJ
• E) 29.3 kJ
1. Consider the following specific heats of metals.
• Metal Specific Heat
• Zinc 0.387 J/(g°C)
• Magnesium 1.02 J/(g°C)
• Iron 0.450 J/(g°C)
• Silver 0.237 J/(g°C)

If the same amount of heat is added to 25.0 g of each of the metals, which are all at the same initial temperature, which metal will have the highest temperature?

• A) Zinc
• B) Magnesium
• C) Iron
• D) Silver
1. Consider the following processes:
 2A –> (1/2)B + C ∆H1 = 5 kJ/mol (3/2)B + 4C –> 2A + C + 3D ∆H2 = –15 kJ/mol E + 4A –> C ∆H3 = 10 kJ/mol

Calculate ∆H for:   C ­­–> E + 3D

• A) 0 kJ/mol
• B) 10 kJ/mol
• C) –10 kJ/mol
• D) –20 kJ/mol
• E) 20 kJ/mol
1. Which of the following does not have a standard enthalpy of formation equal to zero at 25°C and 1.0 atm?
• A) F2(g)
• B) Al(s)
• C) H2O(l)
• D) H2(g)
• E) They all have a standard enthalpy equal to zero.
1. Choose the correct equation for the standard enthalpy of formation of CO(g), where ∆Hf° for CO = –110.5 kJ/mol (gr indicates graphite).
• A) 2C(gr) + O2(g) à 2CO(g),       ∆H° = –110.5 kJ
• B) C(gr) + O(g) à CO(g),           ∆H° = –110.5 kJ
• C) C(gr) + O2(g) à CO(g),       ∆H° = –110.5 kJ
• D) C(gr) + CO2(g) à 2CO(g),     ∆H° = –110.5 kJ
• E) CO(g) à C(gr) + O(g),           ∆H° = –110.5 kJ
1. Consider the reaction:

2ClF3(g) + 2NH3(g) –> N2(g) + 6HF(g) + Cl2(g)

When calculating the ∆H°rxn, why is the ∆Hf° for N2 not important?

• A) Because nitrogen is in its standard elemental state and no energy is needed for this product to exist.
• B) Because any element or compound in the gaseous state requires a negligible amount of energy to exist.
• C) Because the products are not included when calculating ∆H°rxn.
• D) Because nitrogen is in its elemental state and does not contribute to the reaction itself.
• E) Two of the above statements explain why N2 is not important when calculating ∆H°rxn.
1. A 36.2 g piece of metal is heated to 81°C and dropped into a calorimeter containing 50.0 g of water (specific heat capacity of water is 4.18 J/g°C) initially at 21.7°C. The empty calorimeter has a heat capacity of 125 J/K. The final temperature of the water is 29.7°C. Ignoring significant figures, calculate the specific heat of the metal..
• A) 1.439 J/gK
• B) 0.900 J/gK
• C) 0.360 J/gK
• D) 0.968 J/gK
• E) none of these
1. For the complete combustion of 1.000 mole of butane gas at 298 K and 1 atm pressure, ∆H° = -2877 kJ/mol. What will be the heat released when 2.43 g of butane is combusted under these conditions?
• A) –121 kJ
• B) 121 kJ
• C) 68669 kJ
• D) –68669 kJ
• E) none of these
1. Using the following data, calculate the standard heat of formation of the compound ICl in kJ/mol:
 ∆H° (kJ/mol) Cl2(g) –> 2Cl(g) 242.3 I2(g) –> 2I(g) 151.0 ICl(g) –> I(g) + Cl(g) 211.3 I2(s) –> I2(g) 62.8

• A) –211 kJ/mol
• B) –14.6 kJ/mol
• C) 16.8 kJ/mol
• D) 245 kJ/mol
• E) 439 kJ/mol
1. Which has the highest vapor pressure?
• A) 125 mL of water at 283 K
• B) 0.1 mL of water at 288 K
• C) 50 mL of water at 293 K
• D) 5 mL of water at 299 K
1. What quantity of heat is required to change 40.0 g of ice at melting point to liquid water? The heat of fusion of ice is 335 J/g.
• A) 13400 J
• B) 0.119 J
• C) 8.38 J
• D) 375 J
1. What quantity of heat must be removed from 20.0 g of liquid water at 0 °C to completely freeze the water? The heat of fusion of ice is 335 J/g.
• A) 315 J
• B) 16.8 J
• C) 6700 J
• D) 0.0597 J
1. A sample of ice at 0° C absorbs 6030 J of heat energy. How much of the ice can melt? The heat of fusion of ice is 335 J/g.
• A) 6030 g
• B) 2.02 10 6 g
• C) 18.0 g
• D) 5700 g
1. What quantity of heat is required to convert 40.0 g of liquid water at its boiling point to steam? The heat of vaporization of water is 2.26 kJ/g.
• A) 17.7 kJ
• B) 0.0565 kJ
• C) 42.3 kJ
• D) 90.4 kJ
1. A sample of liquid water at 100. °C absorbs 113 kJ of heat energy. How much of the water will be converted to steam? The heat of vaporization of water is 2.26 kJ/g.
• A) 50.0 g
• B) 0.0200 g
• C) 255 g
• D) 113 g
1. A 32.0 g sample of liquid water at 21.0 °C absorbs 2276 J of heat energy. What will be the final temperature of the water? The specific heat of liquid water is 4.184 J/g °
• A) 4.0 0 C
• B) 17.0 0 C
• C) 38.0 0 C
• D) 58.0 0 C
2. Which segment in the following figure indicated by letters corresponds to melting?
• A) AB
• B) BC
• C) CD
• D) DE