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Name ________________________________________ Date ______________ Period _______
HONORS CHEMISTRY – THERMOCHEMISTRY CLASSWORK
DEFINE THE FOLLOWING TERMS (CHAPTER 17):
CALORIE (OR KILOCALORIE)
CALORIMETRY
CHEMICAL POTENTIAL ENERGY
ENDOTHERMIC PROCESS
ENTHALPY
EXOTHERMIC PROCESS
HEAT
HEAT CAPACITY
HEAT OF FUSION
HEAT OF REACTION
HEAT OF VAPORIZATION
JOULE
SPECIFIC HEAT
STANDARD HEAT OF FORMATION
THERMOCHEMICAL EQUATION
THE FOLLOWING TERMS YOU SHOULD ALREADY KNOW FROM PRIOR SCIENCE COURSES:
BOILING
MELTING POINT
BOILING POINT
SUBLIMATION
CONDENSATION
TEMPERATURE
DEPOSITION
VAPOR
EVAPORATION
VAPOR PRESSURE
FREEZING
VAPORIZATION
MELTING
VOLATILE
Name ________________________________________ Date ______________ Period _______
HONORS CHEMISTRY – THERMOCHEMISTRY
Temperature vs. Heat (Previous material)
1. What does motion have to do with matter?
2. How is the Kevin temperature of a substance related to the average kinetic energy of its particles?
3. Describe the difference in kinetic energy for the molecules in hot tea compared to the molecules of
iced tea?
4. How are Temperature and heat related?
5. The diagram to the right represents two different solids of the
material under the same conditions. Block A has less mass
Block B. Answer the following:
a. Compare the temperature of the two blocks.
A
B
same
than
b. Compare the heat energy of the two blocks.
Exothermic vs. Endothermic (Text Section 17.1)
1. Explain what is meant by the terms exothermic and endothermic. Give an example of each.
2. Define the terms system and surroundings. Give an example
3. For the diagram above (Block A & B) suppose Block B is at 75°C and Block A as at 25°C. If they were
pushed in contact with each other:
a. Describe the heat flow that would result.
b. What would be the final temperature of each block compared to their initial temperatures
(increase, decrease, remain the same, etc.).
4. For exothermic reactions:
a. What happens to the temperatures of the environment before and after the reaction?
b. Does energy get absorbed or released into the environment?
c. Where does the energy come from and go to for each reaction?
5. For endothermic reactions:
a. What happens to the temperatures of the environment before and after the reaction?
Name ________________________________________ Date ______________ Period _______
b. Does energy get absorbed or released into the environment?
c. Where does the energy come from and go to for each reaction?
6. In the table below indicate whether the system is exothermic or endothermic.
System
Heat Flow
A snowball in your hand
A fresh cup of coffee
A lit match
A car engine after a long drive
A camp fire
An ice cube on the counter
A snickers bar in a car on a summer day
Your body during PE class
A baked potato
Maggie Moo ice cream in your belly
7. In the space below draw a picture, diagram or any illustration the shows the difference between
Heat and Temperature or between Exothermic and Endothermic.
8. Determine if the following energy changes is Exothermic or Endothermic.
a. The reaction feels cold
d. The reaction releases energy
__________________
__________________
b. The reaction feels warm
__________________
e. Bonds are being broken
__________________
c. The reaction absorbs energy
__________________
f.
Bonds are being formed
_________________
9. Energy always flows from ____________________ to ____________________.
Name ________________________________________ Date ______________ Period _______
10. You are at Ocean city in July. You lay out your towel take your take out your water bottle slip your
shoes off and take off your watch. After about 10 minutes you realize you forgot something back at
the hotel room. You grab your water bottle which is now warm and burn your hand as you grab
your watch. Explain how two objects, sitting in the same place, for the same amount of
time, can be different temperatures.
Energy Unit Conversion (Text Section 17.1)
1. Convert from one unit to the other:
a. 1.69 Joules to calories
b. 423 calories to kilocalories
c. 820.1 J to kilocalories
d. 6.78 kilocalories to kilojoules
Heat Capacity and Specific Heat (Text Section 17.1)
1. Determine the energy required to raise the temperature of 3.21 grams of liquid water 4.0 °C.
The specific heat of water is 4.184 J/g°C.
2. Determine the specific heat of a 150.0 gram object that requires 62.0 Joules of energy to raise
its temperature 12.0 °C.
3. Determine the heat required to convert 62.0 grams of ice at -10.3 °C to 0.0 °C. The specific
heat capacity of ice is 2.1 J/g °C.
4. If 3500 J of energy is absorbed by a 20.0 g piece of lead at 283.0 °C what will the final
temperature of the lead be? The specific heat of lead is 0.128 J/g°C.
Calorimetry (Text Section 17.2)
1. Determine the mass of iron at 85.0 °C needed to add to 54.0 grams of water at 0.0°C to raise the
temperature of the water to 12.5 °C. The specific heat of iron is 0.045 J/g °C.
2. A 15.6 gram piece of sliver at 100.0°C is placed into a sealed vessel with 125 grams of water at
12.5°C. If the mixture has a final temperature of 13.3°C, what is the specific heat of silver?
Name ________________________________________ Date ______________ Period _______
3. HONORS: A sample of cobalt, A, with a mass of 5.00 g, is initially at 25.0 °C. When this sample gains
6.70 J of heat, the temperature rises to 27.9 °C. Another sample of cobalt, B, with a mass of 7.00 g, is
initially at 25.0 °C. If sample B gains 5.00 J of heat, what is the final temperature of sample B. (Hint:
think about the specific heat of both samples.)
4. HONORS: A student places 21.4 grams of copper at 0.0 °C and 13.1 grams of water at 100.0 °C in a
sealed and insulated container. Determine the final temperature of the mixture. The specific heat
capacity of copper is 0.385 J/g°C.
Thermochemical Equations (Text Section 17.2) I recommend you also do Sample Problem 14 & 15
1. Do chemical reactions involve an exchange of energy with the environment (does heat energy
enter or leave during a chemical reaction)?
2. Does energy have an atomic or molecular formula?
3. How would you represent energy flow in this exothermic chemical reaction?
(Example chemical reaction: Zn + 2 HCl → H2 + ZnCl2 )
4. Calculate the amount of heat (in kJ) required to decompoase 2.0 mol hydrogen peroxide in
the following decomposition reaction.
2H2O2(l) → 2H2O(l) + O2(g)
∆H=-98.2 kJ/mole H2O2
Heat of Fusion/Vaporization (Text Section 17.3) I recommend you also do Sample Problem 22 & 27
1. How much heat must be removed to freeze a container of water if the water has a mass of
1.50x102 g?
Name ________________________________________ Date ______________ Period _______
2. The molar heat of vaporization of water is 40.7 kJ/mol. How many kilojoules of heat are
required to vaporize 60.5 g of water at its boiling point of 100°C?
3. How many kilojoules of heat are released when 46.0 g of CaCl2 are dissolved in water?
4. Determine the energy required to:
a. melt 5.62 moles of ice at 0 °C.
c. boil 0.345 moles of water at 100.0 °C.
b. Convert 16.2 grams of ice to liquid
water.
d. Convert 52.6 grams of steam to water.
Hess’s Law & Standard Heats of Formation (Text Section 17.4)
1. Calculate enthalpy change for the following dissociation of sulfuric acid:
()
( )
( )
a.
( )
( )
b.
( )
( )
c.
()
( )
()
( )
()
Calculate the standard heat of reaction for the following: (refer to Table 17.4 in text, pg# 580,
for
values)
2. Thermite reaction: iron(III) oxide is reduced to metallic iron by aluminum (
).
( )
( )
( )
( )
3.
( )
( )
( )