Conceptual Questions
A) The total amount of heat a sample can absorb.
B) The heat required to raise 1 gram of a substance by 1˚C.
C) The same for all substances.
D) The amount of heat needed to melt 1 mole of solid.
B) The heat required to raise 1 gram of a substance by 1˚C.
Which statement explains why water has a much higher specific heat than most substances?
A) It has a low density
B) It has hydrogen bonding between molecules
C) It contains heavy oxygen atoms.
D) It is a liquid at room temperature
B) It has hydrogen bonding between molecules.
In a chemical reaction occurring at constant pressure, the heat change measured in a coffee cup calorimeter corresponds to:
A) Internal energy change (∆E)
B) Enthalpy change (∆H)
C) Gibbs free energy change (∆G)
D) Work done (w)
B) Enthalpy Change (∆H)
102 kJ
Which of the following statements about ∆H vap is true?
A) It represents the heat released when a solid melts.
B) It is typically smaller than ∆H fusion.
C) It represents the energy needed to convert liquid to gas.
D) It has a negative value for vaporization.
C) It represents the energy needed to convert liquid to gas.
The energy needed to convert 1 mole of liquid into gas at its boiling point
A 72.0 g sample of ice is at 0˚C. How much energy is required to convert it to steam at 100.0˚C? ∆Hfusion of water is 6.02 kJ/mol and ∆Hvap of water is 40.7 kJ/mol.
217 kJ
Why does temperature remain constant during melting/freezing and boiling/condensing?
All the heat energy goes into breaking intermolecular forces, not increasing kinetic energy.
Calculate the heat necessary to raise 27.0 g of water from 10.0˚C to 90.0˚C
9.04 kJ
Calculate the heat given off when 159.7g of copper cools from 155.0 ˚C to 23.0 ˚C. The specific heat capacity of copper is 0.385 J/g˚C.
116 kJ
31.2 kJ
Why does condensation release heat?
When gas molecules form intermolecular bonds to become a liquid, energy is released.
How many joules of energy would be required to heat 15.9 g of diamond from 23.6 ˚C to 54.2˚C? (Specific heat capacity of diamond is 0.5091 J/g˚C
248 J
You are given 12.0 g of ice at -5.00 ˚C. How much energy is needed to melt the ice completely to water? C of ice is 2.09 J/g˚C and ∆H vap of water is 6.02 kJ/mol
4.01 kJ
A 10.35 kg block of ice has a temperature of -22.3˚C. The block absorbs 4.696 x 10ˆ6 J of heat. What is the final temperature of the liquid water?
1.76 ˚C
Why is ∆H vap typically larger than ∆H fusion?
Because vaporization completely breaks all intermolecular attractions, while fusion only partially disrupts them.
Calculate the heat given off when 159.7g of copper cools from 155.0 ˚C to 23.0 ˚C. The specific heat capacity of copper is 0.385 J/g˚C.
Lead has a melting point of 327.5 ˚C, its specific heat is 0.128 J/g˚C, and its molar enthalpy of fusion is 4.80 kJ/mol. How much heat, in kJ, will be required to heat a 500.0 g sample of lead from 23.0 ˚C to its melting point and then melt it?
31.1 kJ
7.12 kJ