WPHS Chemistry
Unit 6
States of Matter & Energy
Flow
Bergmann-Sams
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Matter
Chemistry: Unit 6 Outline: States of
Assignment
Podcast 6.1 KMT and Pressure
Worksheet 6.1 KMT and Pressure
Aluminum Can Crush
Podcast 6.2 Vapor Pressure and Phase Diagrams
Demo: Boiling Water with Ice
Worksheet 6.2 Vapor Pressure and Phase Diagrams
Demo: Triple Point of CO2
Podcast 6.3 Energy Transformations
Worksheet 6.3 Energy Transformations
Lab: What factors affect heat flow?
Podcast 6.4 Heat Capacity
Worksheet 6.4 Heat Capacity
Lab: Heat Capacity of a Metal
Podcast 6.5 Energy Stoichiometry
Worksheet 6.5 Energy Stoichiometry
Podcast 6.6 Calorimetry
Worksheet 6.6 Calorimetry
Lab: Heat of Reaction: Magnesium
Podcast 6.7 Phase Change Problems Conceptual
Worksheet 6.7 Phase Change Problems Conceptual
Demo: Burning Paper with Steam
Podcast 6.8 Phase Change Problems Mathematical
Worksheet 6.8 Phase Change Problems Mathematical
Lab: Heat of Fusion of Ice
Unit 6 Vocabulary
Lab Test: Finding ∆H
Unit 6 Exam
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In ClassOnly?






Heat
System
Surroundings
Universe
Law of conservation of energy
Endothermic process
Exothermic process
Calorie
Joule
Heat capacity
Specific heat
Calorimetry
Calorimeter
Enthalpy
Thermochemical equation
Heat of reaction
Heat of combustion
Molar heat of fusion
Molar heat of solidification
Molar heat of vaporization
Molar heat of condensation
Molar heat of solution
Heat of formation
Unit 6: Vocabulary
Kinetic theory
Kinetic energy
Gas pressure
Vacuum
Atmospheric pressure
Barometer
Pascal
Standard atmosphere (atm)
Vaporization
Evaporation
Vapor pressure
Boiling point
Normal boiling point
Melting point
Crystal
Allotropes
Amorphous solids
Phase diagram
Triple point
Sublimation
Thermochemistry
Energy
Chemical potential energy
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Title: Aluminum Can Crush—Take Home Lab
Subject/Concept: Chemistry - Gas Laws
Purpose: The purpose of this activity is to observe the gas laws in action.
Materials:
•1/8 cup water
• small pan of ice water
• 1 empty aluminum soda can
• 1 pair barbeque tongs
• 1 stove
Procedure:
1. Place 1/8 cup of water in an empty aluminum soda can.
2. Heat the can of water on the stove on medium until visible steam escapes from the can
for a period of about 5-10 seconds.
3. With the barbeque tongs, grab the can and invert it into the pan of ice water.
Cautions:
1. Do not boil the can dry! It will melt onto the stove top!
2. Do not touch the aluminum cans with your hands while it is on the stove. It is
very hot.
3. Be careful not to spill or splash any of the boiling water in the can as you invert it.
Questions:
1. What gas replaces the air while the water is boiling?
2. What is the chemical formula for the gas in question #1?
3. Explain how each of the following contributes to the results:
a) the temperature of the gas cools when the can is removed from the heat
b) the gas inside the can condenses to form water droplets
4. What is the difference between an implosion and an explosion?
5. If cans are roughly 8” in circumference and 4.75” tall, how many pounds of pressure
does the atmosphere place on the outside of the can? (1 atm = 14.7psi )
For Credit:
To receive credit, your parent or guardian must write a short note confirming that you
performed the experiment for them and explained the results to their satisfaction using the
concept of gas laws. Attach your note to the back of this sheet.
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Lab: What factors affect heat flow?
Pre-Experiment Discussion
Chemical and physical changes are always accompanied by a change in energy. Most commonly this
energy change is observed as a flow of heat energy.
Previous Knowledge:
 Understand basics of heat transfer
 Knowledge of types of energy
Questions:
1. Two objects with different temperatures are touching. Describe the heat flow.
 Which direction does the heat flow? Why?

How will the temperatures of each object change?

When will the heat flow cease?

What will the final temperatures be?
2. Discuss with your group the Question of the Day. Present a summary of your ideas to the rest of the
class.
Materials:
beaker, 400-mL
Bunsen burner/hotplate
wire gauze
ring stand
test tube tongs
electronic balance
4 materials in small pieces
test tube, large
thermometer
graduated cylinder, 100-mL
Styrofoam cups, 2, with lid
Safety Alert:

All materials being used around the Bunsen burner should be assumed to be hot. Metal
and glass looks the same hot as it does cold. Use heat resistant gloves to handle hot objects.
Wear Safety Goggles the entire time!
Procedure:
1. Set up your Bunsen burner, wire gauze, and ring stand. Fill your 400-mL beaker about
4/5 full of tap water. Set it on your wire gauze. Light the Bunsen burner so the water can
be heating as you prepare the sample of metal.
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2. Using the electronic balance, obtain about _____ g of the __________. Record the exact
mass on your data table. Put an X in the box for the sample that you will be studying.**
other ideas: sand, soil, antifreeze, water,
Masses
40 g
60 g
80 g
Substances
Aluminum
Copper
Zinc
Glass
3. Add the substance to your large test tube. Put your thermometer in the middle of the
pieces so it can track the temperature of the sample. It might be easier to add some
pieces, then insert the thermometer, and then add the rest of the metal pieces around it.
4. Gently place the test tube in the warm water. Be sure that the entire sample of the
substance is below the water line. Answer In-Lab Question #1.
5. Prepare your calorimeter by stacking two Styrofoam cups together.
6. Measure exactly 100 mL of tap water into the calorimeter.
7. Insert your second thermometer into the opening in the lid of your calorimeter. Record
the temperature of your water in the data table.
8. Allow the metal to reach a temperature between 85 and 90 °C. Measure and record the
exact temperature of the metal in the test tube in your data table.
9. Answer In-Lab Question #2. Work quickly but carefully: remove the lid of the
calorimeter and empty the hot metal from the test tube into the calorimeter. Replace the
calorimeter lid and thermometer. Set the test tube aside.
10. Answer In-Lab Question #3. Gently swirl the mixture of metal and water in your
calorimeter. Watch the reading on the thermometer. Record the highest temperature
reached for the mixture in your data table. Answer In-Lab Question #4.
11. To cleanup, remove the thermometer from the calorimeter. With the lid on your
calorimeter to catch the metal pieces, pour the water into the sink. Put the wet metal
caught in the lid in the labeled container to dry.
12. Repeat Steps 1-11 two more times.
In-Lab Questions:
Discuss and answer questions with your partner at the appropriate times noted in the procedure.
1. Is the temperature of the metal before heating important? Explain.
2. Why would the hot metal need to be transferred to the water quickly?
3. What is the purpose of the swirling?
4. What would happen to the temperature of the mixture if it was allowed to sit longer?
Explain.
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Data Table
Pair Data Collected
Type of Substance
Trial 1
Trial 2
Mass of substance
Temperature of water
Temperature of substance
Temperature of mixture
Trial 3
g
g
g
°C
°C
°C
°C
°C
°C
°C
°C
°C
Post-Experiment Questions:
Discuss and answer questions with your partner at the appropriate times noted in the procedure.
After you have finished, compare your answers with the other members of your larger group and
discuss any differences.
1. Which part of the mixture, the substance or the water, was releasing heat? Which was
absorbing heat? How do you know?
2. What can you say about the final temperature of the objects?
3. Calculate the change in temperature of the water for each trial and then calculate the
average of these values.
4. Calculate the change in temperature of the metal for each trial and then calculate the
average of these values.
5. Compare the change in temperatures of the three metal samples. Describe any trend you
find.
6. Propose a hypothesis for the pattern you have observed.
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7. Compare the change in temperature of the water to the change in temperature of the
metal. Describe any trend you find.
8. Propose a hypothesis for the pattern you have observed.
9. Create one line graph of mass of metal vs. the inverse of the average change in
temperature of the metal (mass vs. 1/∆T) for each of the four metals. Make each of the
four lines a different color. Add a line of best-fit for each metal and display the equation
on the chart.
10. Discuss with your group the meaning of the slope. Summarize your ideas here.
11. Which metal has the highest slope? Which has the lowest slope?
12. What do these differences suggest about how these metals transfer heat?
13. Based on the data from this experiment, summarize the factors that affect heat transfer.
14. Brainstorm properties of the different materials that might account for their different heat
flow behaviors.
15. Using grammatically correct sentences compare the heat transfer ability of each material
tested.
16. If a similar experiment was done using 100 mL of water at 20°C with 100 g of metal at
80°C, what would you expect the approximate final temperature to be? Explain.
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17. Assuming an island and inland areas or exposed to the same amount of heat energy, why
would the island have less drastic temperature changes than inland area?
18. How would you determine Styrofoam’s ability to transfer heat? What difference, if any,
would you find in its behavior compared to metal? What is a common use of Styrofoam
that capitalizes on this idea?
19. If 5 g samples of glass and copper are placed on the same hot plate and allowed to heat
for one minute, how would their final temperatures compare? Explain.
20. Reflecting on what you have learned throughout the experiment, summarize what you
have learned about heat transfer.
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WPHS Chemistry
Identification of an Unknown Metal
In this lab we will be using lab techniques and basic chemical concepts to identify an
unknown metal. Every metal has a unique set of properties. We will be using specific
heat (also known as "heat capacity" or "specific heat capacity"). Your periodic tables
have adequate listings for the purpose of this experiment, but several other sources also
have listings of these values for pure materials: check the indices of the CRC Handbook,
the Merck Index, or the Exploring the Elements books on the shelves around the room.
NOTES: Goggles are necessary, and long hair must be tied back . Record all data to the
correct decimal place in your lab handbook. Since we are using digital thermometers,
feel free to stir with the thermometer.
PROCEDURE - Specific Heat
To determine the specific heat of a metal sample, we will use a calorimeter, and the
concept that in a closed system, heat lost by a hot object is gained by a cooler one.
mCp∆T (cold) = –mCp∆T (hot)
To determine the initial (high) temperature of your metal sample, suspend it in a beaker
of boiling water and keep it there until boiling has proceeded steadily for about two
minutes. Record the temperature of the boiling water with a thermometer in a rubber
stopper (it should not touch the beaker glass). This is the same as the initial metal
temperature.
While the water/metal mixture is boiling, record the mass of the empty calorimeter cup.
Next, add just enough water to cover your piece of metal (estimate), and determine the
combined mass. Keep track of the temperature of this water with a second thermometer.
Record the temperature of the cold water and of the metal just before combining them.
Immerse the metal sample in the cold water and record the final temperature of the
mixture. It should change quickly at first, then level off, then cool back down slowly.
Record the level part.
NOTE: Perform one density trial then one specific heat trial, then repeat the whole thing
two more times. Make sure Everything is dry for each trial.
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Data Table
Trial 1
Trial 2
Trial 3
initial temperature of metal __________ __________ __________ (C)
temperature of cold water __________ __________ __________ (C)
final temperature of mixture __________ __________ __________ (C)
specific heat of water
4.18
4.18
4.18
(J/gC)
mass of cold water (use Vol) __________ __________ __________ (g)
change in water temperature __________ __________ __________ (C)
change in metal temperature __________ __________ __________ (C)
specific heat of metal
__________ __________ __________ (J/gC)
Show a complete sample calculation for heat capacity and then determine which metal
you have using the table below
Metal
Specific
Heat (J/g
ºC)
Aluminum
0.91
Cast Iron
0.46
Iron
0.46
Molybedenum 0.25
Tin
0.21
Metal
Antimony
Copper
Lead
Nickel
Titanium
Specific
Heat (J/g
ºC)
0.21
0.39
0.13
0.54
0.54
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Metal
Specific
Heat (J/g
ºC)
Carbon Steel 0.49
Gold
0.13
Magnesium 1.05
Sliver
0.23
Zinc
0.39
HEAT OF REACTION OF MAGNESIUM
Description
The mass of a measured length of polished magnesium ribbon is determined. A small length of ribbon is
cut and its length measured. From the length of the small piece, its mass is calculated. A known volume of
acid is placed in a calorimeter. The temperature is measured. The magnesium is added to the calorimeter,
and a reaction takes place. The temperature is measured once reaction ceases. From the mass of ribbon
reacted, the temperature increase, and the volume of acid used, the heat of formation of Mg2+(aq) is
determined.
Hazards
Magnesium burns in a very exothermic reaction. Never look directly at burning magnesium. Cuts when
polishing the magnesium metal are possible. The dilute acid is corrosive.
Precautions
Keep magnesium away from ignition sources such as open flames. Wear gloves when polishing the
magnesium metal ribbon. Wear eye protection; wash acid spills immediately.
Procedure
Demonstration:
1.
Place 100 mL of 3 M hydrochloric acid in a 250-mL beaker. Fold a 10-cm length of magnesium
ribbon, place in the solution, and note evidence for reaction.
Experiment:
1.
2.
3.
4.
Carefully measure and record the length of one of the cut pieces of polished magnesium ribbon.
Obtain the mass of a 1.00 meter length of polished magnesium ribbon from your instructor.
Crumple the piece of magnesium into a small ball.
Pour 50-60 mL of 3 M hydrochloric acid into a 100-mL cylinder. Measure and record the exact
volume to the nearest 0.2 mL.
Assemble a calorimeter.
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5.
6.
7.
8.
A calorimeter assembled from three foam cups works well. Two nested cups hold the fluid. One
fourth of the third cup is removed at the lip. A hole is made in the bottom for the thermometer; a
hole is made in the side near the upper edge through which additions can be made. When
inverted, this piece serves as a calorimeter cover.
Pour the measured volume of acid into the inner calorimeter cup. Cover. Insert the thermometer.
Read and record the temperature to the nearest 0.1 °C at regular intervals until it becomes
constant.
Add the crumpled magnesium to the acid solution. Swirl gently. Note the temperature.
Record the maximum temperature reached by the hydrochloric acid solution.
Data
length magnesium ribbon (cm)
mass Mg/ m (provided by
instructor)
mass Mg used
mol Mg used
volume HCl (mL)
initial temperature, °C
maximum temperature, °C
temperature rise, °C
energy released (J)
energy released/ mol Mg (kJ/mol)
Percent Error
Accepted value = -462.0 kJ/mol at 25 °C for 1 M H+
1.
2.
3.
Write an equation for the reaction of magnesium metal with hydrochloric acid. Include the heat of
reaction calculated above.
State the relationship between the heat of reaction for this reaction and the heat of formation of
aqueous magnesium ion.
The accepted heat of formation of Mg2+(aq) is -462.0 kJ/mol (25 °C, 1 M). Find the percent error
for the experiment
- 14 -
Heat of Fusion of Ice
Name
INTRODUCTION:
The amount of energy required to convert a solid to a liquid, at constant pressure and
temperature, is called the heat of fusion of the substance. In this experiment, the heat of
fusion of ice will be determined.
The ice will be melted by placing it in a known volume of hot water contained in a plastic
cup. The system will be left undisturbed until all the ice has melted. The amount of heat
lost by the hot water in this process can be calculated according to the following
equation.
Procedure
1. Get 100.0 mL of water at the highest temperature you can out of the tap.
2. Place the water in a Styrofoam cup.
3. Measure the temperature of the water.
4. Simultaneously, the other partner should be getting a handful of ice cubes that
need to be dried.
5. Place the ice cubes in the cup of hot water and wait until the ice cubes have
completely melted.
6. Measure the final temperature.
7. Measure the total volume of the cool water.
8. Place all data in the table below.
9. For the masses: Assume that the density of water is 1.0 g/mL.
DATA TABLE:
Initial mass of hot water.
g
Initial temperature of hot water.
ºC
Final temperature of water and melted
ice.
ºC
Final mass of water and melted ice.
Change in Temp of Hot Water
Change in Temp of Ice
mass of ice added
g
ºC
ºC
g
- 15 -
Calculations
1.
Assuming that the heat gained by the ice (Qice) is equal to the heat lost by the warm water
(Qwarm water) you can solve for the value of ΔH for ice.
mice cΔT + ΔHfusmice = mwarm water cΔT
Where the left side of the equation represents the ice melting and then warming up to the final
temperature and the right side represents the energy lost by the warm water.
Now solve for ΔHfus
.
Questions
1.
The accepted value for the heat of fusion of water (ice) is 334 joules/g. What is your percentage of
error?
2.
What is the most important source of error?
3.
What happened to the temperature of the ice as it was melting?
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Worksheet 6.1: KMT and Pressure
Name _____________________
1. Use kinetic theory to explain what causes gas pressure.
2. Convert the following Pressures:
a. 600 mm Hg into atm
b. 190 kPa into atm
c. 2.3 atm into mm Hg
d. 1.5 atm into kPa
3. What is the equation that relates kinetic energy to Kelvin Temperature
4. What is the temperature at absolute zero
5. Use kinetic theory to explain the difference between evaporation and boiling of a
liquid.
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6. Use the chart to answer the following questions.
a. What is the vapor pressure of ether at 40ºC?
b. What is normal boiling point of water?
c. What is the normal boiling point of benzene?
7. How is the average kinetic energy of water molecules affected when you pour hot
water from a kettle into cups at the same temperature as the water?
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Worksheet 6.2 Vapor Pressure and Phase Diagrams Name __________________
1. What is the relationship between atmospheric pressure and altitude? What effect does
this have on the boiling point of water?
2. Define vapor pressure.
3. What is the full definition of the boiling point of a liquid?
4. Explain why the boiling point of a liquid varies with atmospheric pressure.
5. Using the diagram above: What would happen to the CO2 if you:
a. Heat it up from -60º to 30º C at 6.0 atm?
b. Increase the pressure from 2.0 atm to 60 atm at a temperature of 0ºC?
c. Decrease the pressure from 80.0 atm to 1 atm at a temperature of 25ºC?
d. Increase the temperature from -80ºC to +80ºC at a pressure of 1.0 atm?
e. Increase the temperature from -60ºC to 20ºC at a pressure of 70 atm?
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WS 6.3: Energy Transformations
Name ________________
1. Which has more heat: Rampart Reservoir or a cup of boiling water? Explain
your answer.
2. What always happens when two object of different temperatures come into
contact? Give an example from your daily life.
3. Classify these process as exothermic or endothermic:
 Condensing steam

Evaporating alcohol

Burning alcohol

Baking a potato
4. Energy always flows from _______ to _____________.
5. When you open your door in the winter and your mom says: “Quit letting the
cold in the house,” what should you politely tell her?
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WS 6.4 Heat Capacity
Substance Heat
Capacity
(J/gºC)
Water
4.18
Substance Heat
Capacity
(J/gºC)
Grain
2.4
Alcohol
Ice
2.1
Steam
1.7
Chloroform 0.96
Aluminum 0.90
Glass
0.50
Iron
0.46
Silver
0.24
Mercury
0.14
Lithium
0.14
Copper
0.39
Uranium
0.12
Gold
0.13
Use the table above to answer the following questions
1. A 75.0g sample of a metal at 98.0º C is dropped into a container of 350. g of
water at 24.0ºC, The final temperature is 24.5 ºC. What is the specific heat of the
metal?
2. A 55.0 g sample of a metal at 90.0º C is dropped into a container of 250 g of
water at 22.0ºC, The final temperature is 23.4 ºC. What is the metal?
3. A 200.0 g sample of a metal at 92.0º C is dropped into a container of 100 g of
water at 26.0ºC, The final temperature is 45.6 ºC. What is the metal? What is the
percentage error?
4. A 45.2 g sample of Silver at 101.0 ºC is dropped into 300.0 g of water at 30.6 ºC.
What is the final temperature of the water?
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WS 6.5: Energy Stoichiometry
1.
Name __________________
Nitrogen reacts with hydrogen according to the following equation.
N2 + H2  NH3 + 231 kJ
How much energy will be released when 13.0-g of nitrogen reacts?
How much energy will be released when 2.5L of hydrogen reacts?
2. Carbon disulfide is an important industrial solvent. It is prepared by the reaction
of coke with sulfur dioxide:
5C(s)+2SO2(g)  CS2(s)
+4CO(g)
∆H = 23.5-kJ
How much energy will be required when 32.4-g of Carbon disulfide is produced?
How much energy will be required to react 21.1-L of sulfur dioxide?
3. Silver nitrate reacts with calcium to make calcium nitrate and silver. The reaction
is exothermic and produces 18.7 kJ/mol of energy.
AgNO3+ Ca Ca(NO3)2 + Ag
How much energy will be produced when 12.3-g of silver nitrate react?
How many grams of silver nitrate reacted if 13.7 kJ of energy is released?
- 22 -
4. Magnesium reacts with hydrochloric acid in a single replacement reaction. The
value of ∆H =-221 kJ/mol. (You will need to write the balanced chemical
equation.
How many grams of magnesium will be made if 365-kJ is released?
How much energy will be released if 9.45-g of magnesium reacts?
5. Ammonium sulfate reacts with barium hydroxide endothermically. ∆H = + 127
kJ/mol. (You will need to write the balanced chemical equation.)
How much energy will be required to react completely 34.5-g of barium
hydroxide?
If 395-kJ of energy is absorbed, how many grams of ppt will be formed?
- 23 -
WS 6.6 Calorimetry
1. When 12.3-g of magnesium reacts with 1000.0g of hydrochloric acid it raises
the temperature from 22.5˚C to 48.3˚C. What is the value of ∆H?
2.
When 13.5-g of ammonium nitrate is dissolved in water it cools 1000.0-g of
water from 32.3˚C to 29.5˚C. What is the value of ∆H?
3. The combustion of 25.6-g of propane, C3H8 raises the temperature of 1000-g of
water by 16.8˚C. What is the value of ∆H?
4. The single replacement reaction:
Li + Al2(SO4)3  Li2SO4 + Al
has a ∆H of –229.7 kJ/mol. If 12.1-g of Li reacts, how much will the temperature of
1000.0-g of water rise? (Hint: balance the reaction and then use Q=mc∆T and solve
for ∆T)
5. For the dissolving of sulfuric acid, H2SO4 the value of ∆H=-236 kJ/mol. If 2.54-g
of sulfuric acid dissolves in 100.0-g of 20.0˚C water, what will be the final
temperature of the water?
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WS 6.7: Phase Change Problems Conceptual
Name____________________
C
1. In the diagram above, label all of the states of matter.
2. Assuming the above diagram is water, label the temperatures of the flat portions
of the diagram above.
3. In the first rising portion of the graph (a), describe what is happening as energy is
added. Discuss this in terms of the kinetic molecular theory.
4. In the 1st flat portion of the graph (b), describe what is happening as energy is
added. Discuss this in terms of bonding forces of attraction.
5. In the 2nd rising portion of the graph (c), describe what is happening as energy is
added. Discuss this in terms of the kinetic molecular theory.
6. In the 2nd flat portion of the graph (d), describe what is happening as energy is
added. Discuss this in terms of bonding forces of attraction.
7. In the 3rd rising portion of the graph (e), describe what is happening as energy is
added. Discuss this in terms of the kinetic molecular theory
- 25 -
8. What is the formula used for calculating the heat involved in a phase change?
9. During a phase change, how much does the temperature change? How much does
the KE change? Does the potential energy change?
10. What is the formula used to calculate the heat required to warm or cool one phase
of matter?
11. When you heat one phase of matter, how do you know that the KE changes?
12. What is meant by the negative sign in an answer like "-46.8 kJ"? When would
you use a positive sign?
- 26 -
WS 6.8: Phase Change Problems Mathematical Name____________________
m.p.
(ºC)
0.00
∆Hfus
(kJ/g)
0.333
b.p.
ºC
100.00
∆Hvap
(kJ/g)
2.26
Grain
Alcohol
-98
0.0987
64
1.10
Benzene
5.0
0.1265
80
0.394
Substance
H2O
Specific Heat –C(J/g ºC)
Ice: 2.09
Water: 4.18
Steam: 1.7
Solid: 1.2
Liquid: 2.4
Gas: 1.9
Solid: 0.55
Liquid: 0.96
Gas: 1.09
1. If you must add 25 kJ to raise the temperature of an ice cube from -15°C to -10°C,
is this an endothermic or an exothermic process.
2. How much heat is required to raise 40 grams of water from 30°C to 70°C?
3. How much water can be raised from 25°C (room temperature) to 37°C body
temperature by adding the 2,000 kJ in a Snickers Bar?
4. How much heat does it take to melt 65 grams of ice at 0°C?
- 27 -
5. Calculate the amount of energy required to change 100 grams of solid ice at 0°C
to gaseous steam at 100°C. (Be aware of units: kJ or J?) How many steps does
this take?
6. Calculate the amount of energy released by cooling 59 grams of liquid water from
+25°C to ice at -25°C. (Be aware of units: kJ or J?) How many steps does this
take?
7. How much heat would it take to raise 5 grams of H2O from -50°C to +200°C?
How many steps does this take?
8. How much energy is required to bring 45 g of benzene from - 10˚C to 70˚C?
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