Name:__________________________________________ Date:____________________ Period:_______
Hot and Cold Pack Unit (chapter 17)
Thermochemistry
This chapter has the chemistry you need to know to determine, from the stand point of a product
developer, which reaction would produce the best hot or cold pack for the cost.
Introduction:
You are high in the mountains on a camping trip, and it begins to snow. Your hands are cold, and
you reach for a hand warmer package. You shake the content as instructed, and place the hot pack
inside your gloves. How does it work?
You’re running in a track meet and injure your ankle. The athletic trainer snaps a cold pack and
places it on your ankle. How does it work?
This unit will explain both the physical and chemical processes and the energies associated with
them.
Name:__________________________________________ Date:____________________ Period:_______
POGIL – Heat & Temperature
Why?
Lava flowing out of an erupting volcano is very hot. Its temperature ranges from 550˚C
to 1400˚C. As lava flows down the side of a volcano, it loses heat and begins to cool
slowly. In some instances, the lava may flow into the ocean, where it cools more rapidly. In this POGIL, you
will learn more about heat flow and why some substances cool down or heat up more quickly than other. You
will also be able to understand the relationship and distinguish between heat and temperature.
Model 1: Lab Introductory Activity
Experiment Overview:
The purpose of this advanced inquiry lab is to design an effective hand warmer that is inexpensive, nontoxic,
and safe for the environment. The investigation begins with an introductory activity to become familiar with
the principles of calorimetry and heat of solution calculations (model 2). The results provide a model for the
guided-inquiry challenge, which is to design an optimum hand warmer for consumer applications. Working in
groups of four, each student group will be provided six different solids, along with their costs and individual
Material Safety Data Sheets (MSDS). Determine the heat of solution for each solid and analyze the cost and
safety information to propose a design for the best all-around hand warmer.
Information:
Hot and cold packs are frequently used by athletes and trainers to treat minor
injuries, such as, inflammations, sprains, muscle spasms, head-aches, etc. Many
commercial hand warmers consist of a plastic package containing a solid and an
inner pouch filled with water. When the pack is activated, the solid dissolves
in water and produces a large temperature change. This generates a chemical
reaction. Depending on the type of salt, this reaction can either release or absorb
heat energy. A hot pack is produced if an exothermic reaction occurs as the salt
and water mix and heat energy is released because this process will raise the
temperature of the contents in the pack. A cold pack is produced if an
endothermic reaction occurs as the salt and water mix and heat energy is
absorbed because this process will lower the temperature of the contents in the pack. The amount of heat that
is released or absorbed by the packs depends on the concentration of water and salts.
The energy or enthalpy change associated with the process of a solute dissolving in a solvent is called the heat of
solution (ΔHsoln). At constant pressure, this enthalpy change, ΔHsoln, is equal in magnitude to the heat (q) loss
or gain to the surroundings.
Heats of solution and other enthalpy changes are generally measured in an insulated vessel called a
calorimeter that reduces or prevents heat loss to the atmosphere outside the reaction vessel. The process of a
solute dissolving in water may either release heat into the resulting aqueous solution or absorb heat from the
solution, but the amount of heat exchanged between the calorimeter and the outside surroundings should be
minimal. When using a calorimeter, the reagents being studied are mixed directly in the calorimeter and the
temperature is recorded both before and after the reaction has occurred. The amount of heat transfer (q) may
be calculated using the heat energy equation:
Equation 1:
q = m × C × ΔT
Name:__________________________________________ Date:____________________ Period:_______
where m is the total mass of the solution (solute plus solvent), C is the specific heat of the solution, and ΔT is
the observed temperature change (final temperature – initial temperature). The specific heat of the solution is
the amount of heat required to raise the temperature of 1 gram of the solution by 1˚Celsius and is
generally assumed to be the same as that of water, namely, 4.18 J/g∙°C.
When measuring the heat transfer for an exothermic heat of solution using a calorimeter, most of the heat
released is absorbed by the aqueous solution (qaq). A small amount of the heat will be absorbed by the
calorimeter itself (qcal). The overall heat transfer (qsoln) for the reaction (the system) then becomes:
Equation 2:
qsoln = –(qaq + qcal)
In order to determine the correction factor qcal for heat of solution calculations, the heat capacity of the
calorimeter, also called the calorimeter constant, must be determined experimentally. The calorimeter constant
has units J/oC. This calibration experiment is done by mixing equal volumes of hot and cool water in the
calorimeter and measuring the temperature after 20 seconds. The resulting value is assumed to be the
instantaneous mixing temperature, Tmix. The average temperature Tavg of the initial hot (TH) and cool water (TC)
is also calculated: Tavg = (TH + TC)/2. The difference between Tavg and Tmix is due to the heat lost by the water and
absorbed by the calorimeter. The heat lost by the water, qwater, is:
Equation 3: qwater = (mass of water) × (specific heat of water) × (Tmix – Tavg)
where the mass is the total mass of hot and cool water.
The heat gained by the calorimeter, qcalor, is equal to that lost by the water, but opposite in sign, (qwater = -qcal)
The calorimeter constant, Ccal, is calculated as follows:
Equation 4:
where Tinitial is the initial temperature of the calorimeter containing cool water.
To calculate the correction factor qcal for use in Equation 2 above—to determine the heat of solution or heat of
reaction for any system—the calorimeter constant is multiplied by the change in temperature of that solution.
qcal = ΔT (°C) × Ccal (J/°C)
Safety Precautions:
Ammonium nitrate is a strong oxidizer and may explode if heated under confinement. It is also slightly toxic by ingestion
and a body tissue irritant. Calcium chloride is slightly toxic. Lithium chloride is moderately toxic by ingestion.
Magnesium sulfate is a body tissue irritant. Sodium acetate is a body tissue and respiratory tract irritant. Avoid contact
of all chemicals with eyes and skin. Wear chemical splash goggles, chemical-resistant gloves, and a chemical-resistant
apron. Wash hands thoroughly with soap and water before leaving the laboratory.
Name:__________________________________________ Date:____________________ Period:_______
Part A. Heat Capacity of the Calorimeter
1. Working in you groups, set up a calorimeter consisting of two nested
polystyrene cups in a ring clamp attached to a support stand.
2. Place a magnetic stirrer below the calorimeter, then lower the ring
clamp until the bottom of the cup just sits on the surface of the
magnetic stirrer (see Figure 1).
3. Measure 100.0 mL of distilled water in a 100-mL graduated cylinder and transfer the water into the
calorimeter.
4. Add a magnetic stirring bar to the calorimeter, and set the bar spinning slowly. If a magnetic stirrer is
not available, use a stirring rod. Do not remove the stirring rod from the calorimeter.
5. Measure and record the initial temperature of the water.
6. Heat approximately 125 mL of distilled water to about 70 °C in a 250-mL beaker.
7. Using heat-resistant gloves, measure 100.0 mL of the 70 °C distilled water in a 100-mL graduated
cylinder.
8. Measure and record the temperature of the hot water.
9. Immediately pour the hot water into the room temperature water in the calorimeter.
10. Insert the thermometer, and stir the water.
11. Record the mixing temperature Tmix after 20 seconds.
12. Empty the calorimeter and dry the inside.
Part A. Heat Capacity of the Calorimeter
Volume of Deionized Water, Cold
Temperature, Cold Water(Tinitia)
Volume of Deionized Water, Hot
Temperature, Hot water
Final Temperature (Tmix)
13. Calculate the calorimeter constant, Ccal, using Tmix (re-read the background section) and Equations 3
and 4 from the Background section.
Name:__________________________________________ Date:____________________ Period:_______
Part B. Calorimetry Procedure
Working in pairs, examine the heat energy change for the following solution.
MgSO4(s) + H2O(l) → Mg2+(aq) + SO42–(aq)
1. Measure 45.0 mL of distilled or deionized water in a 100-mL graduated cylinder and transfer to the
calorimeter.
2. Measure and record the initial temperature of the water.
3. Measure 5.00 g of anhydrous magnesium sulfate in a weighing dish.
4. Put a magnetic stir bar or stirring rod into the calorimeter and slowly stir the water.
5. Quickly add the 5.00 g of anhydrous magnesium sulfate to the calorimeter and insert the thermometer.
6. Monitor the temperature and record the highest or lowest temperature reading.
Part B. Calorimetry Anhydrous
Volume of Deionized Water
Mass of MgSO4
Initial Temperature
Final Temperature
7. Calculate the molar heat of solution for magnesium sulfate. Include the correction due to the heat
capacity of the calorimeter.
Where qcal = ΔT (°C) × Ccal (J/°C)
and
qaq = msoln x ΔT (°C) × C (J/g°C)
Questions about q
Name:__________________________________________ Date:____________________ Period:_______
NH4Cl
MgSO4
Surrounding
Surrounding
System is
the Reaction
+q




System is
the Reaction
-q
Endothermic Reaction: system
gains heat while the
surrounding cools down.
Endothermic processes result
in heat leaving the
surroundings and entering
the system
Mathematically, heat entering
a system has a positive sign
(+q)
Endo takes in heat, so there is
more heat in the system after
the reaction takes place, thus
q is positive.




Exothermic Reaction: system
loses heat as surrounding
heats up.
Exothermic processes result
in heat leaving the system
and entering the
surroundings.
Mathematically, heat
entering the surroundings
has a negative sign (-q)
Exo releases heat, so there is
less heat in the system after
the reaction takes place, thus
q is negative.
You were measuring
in the lab. So with
the MgSO4, the surrounding heated up (temperature
increased). With the NH4Cl, the surroundings cooled
down (temperature decreased).
Model 2: Endothermic & Exothermic Reactions
Thermochemistry is the study of energy changes and transfers that occur during chemical reactions.
Name:__________________________________________ Date:____________________ Period:_______
Energy is the capacity for doing work or supplying heat. Heat deals with energy flow. Heat, represented by q,
is a form of energy. Temperature is a ratio of energy per molecule. Heat and temperature are not the same,
but there is a connection. We use temperature to measure of how hot or cold something is. It is measured in
degrees Celsius (°C) using a thermometer. The two are very different.
Thermochemistry focuses on the study of heat transfer (or heat flow) between the system and the
surroundings. A system is a portion of the universe in which you have an interest, it simply indicates what
you have your attention focused on for the purpose of the problem at hand. Surroundings, in contrast,
represent everything else in the universe that is not in system.
Heat transfers from one object to another naturally as a result of the temperature difference between them.
Heat always flows from warmer objects to colder ones. Heat will continue to flow until the temperature
difference between them is zero.
During chemical reactions, energy (in the form of heat) can either be absorbed by the reaction called
endothermic reaction. Endothermic processes result in heat leaving the surroundings and entering the system.
A cold pack oes not really produce cold-it absorbs the heat from your body (or whatever it touches), and as the
heat leaves your body (the surroundings), you feel cold. Heat can also be released by the reaction called an
exothermic reaction. Exothermic processes result in heat leaving the system and entering the surroundings.
Mathematically, heat entering a system (endothermic) has a positive sign (+q), and heat entering the
surroundings (exothermic) has a negative sign (-q). Energy stored in a compound is really stored in its
chemical bonds – this is referred to as chemical potential energy. Sometimes, energy is required to break
bonds; sometimes, it is required to form them.
The law of conservation of energy states that energy can be neither created nor destroyed; therefore, the sum
of the total heat in the system and the surroundings must remain the same. This means that as a certain
amount of heat leaves the system, the same amount must enter the surroundings (and vice versa).
Critical Thinking Question
1. Which figure to the right, a or b, is an endothermic process? _______. Explain why.
2. Which figure to the right, a or b is an
exothermic process? ________. Explain
why.
3. In the first column of the TEJ, draw an
arrow to indicate the direction of heat
flow. Then in the second column
identify the system and surroundings and in the third column identify whether the system is
undergoing and exothermic or endothermic process.
Name:__________________________________________ Date:____________________ Period:_______
Indicate the direction
of heat flow with an
arrow
Identify the System and
Surroundings
Is the system undergoing
an exothermic or
endothermic process?
System - Snowball
Example 
Endothermic
Surroundings - Hands
Snowball melting
System Surroundings Match burning
System Surroundings Egg cooking on sidewalk
System Surroundings Ice melting in water
System Surroundings Cold pack on arm
System Surroundings Sublimation
Model 3: Calorimetry and Heat Capacity
Name:__________________________________________ Date:____________________ Period:_______
We saw in model 2 that Heat is the flow of energy from an object with higher temperature to one with a lower
temperature. Enthalpy (∆H) is a quantity that takes into account the internal energy of a system, as well as the
pressure and volume of the system. As long as the system pressure does not change, the enthalpy is equal to
the energy flow (or heat,q) of the system. All the work done in the lab will take place in an “open system”meaning that it’s open to the atmosphere pressure around us. Since there is very little pressure changes in the
atmosphere, we can consider enthalpy (∆H) and heat(q) are essentially the same thing. Scientist often use the
term enthalpy, so we will too.
The two common units used for measuring Energy are the calorie and the SI unit joule. A calorie is defined as
the amount of energy needed to heat 1 gram of water by 1˚C. This applies only to a calorie when written with
a small “c” in its name. When written with a capital “C”, as in Calorie¸ it refers to the more commonly known
dietary calorie. Hence:
1 kilocalorie (kcal) = 1000 calories (cal) = 1 Calorie (Cal)
4.184 J = 1 cal
Temperature is measured in degrees. Temperature is most commonly measured in degrees Celsius or Kelvins,
and less commonly in degrees Fahrenheit. Interconversions between temperature scales are common and can
be determined using the equations below:
K = C + 273.15
F = 1.8C + 32
Where K = Kelvins, C = degrees Celsius and F = degrees Fahrenheit
Critical Thinking Questions
4. Using the conversion factors above, make the following conversions using dimensional analysis. Show
all work
a. How many joules are in 12.7 cal?
b. How many calories are in 3.97 kJ?
c. How many calories are in 3.97 kJ?
d. Convert 444 cal to joule
e. Convert 8.50 x 102 cal to Calories
f.
Which substance boils at a higher temperature, ethanol (BP = 351.2 K) or methyl mercaptan (BP =
253.4 0F)? Show your work.
Information:
Name:__________________________________________ Date:____________________ Period:_______
Why does it take some pans a long time to heat on the stove, when others get hot very quickly? Some
molecules or atoms can absorb a lot of energy within their own structures before added energy causes them to
move faster. When they do reach the point where they begin to move enough faster for the thermometer to
detect a change, their temperature will rise to reflect this faster motion.
A substance’s ability to absorb energy before it changes temperature observably is called its heat capacity.
Heat capacity is the amount of energy a substance can absorb before noticeable increase in temperature.
Scientists have defined specific heat capacity (C or Cp) as the amount of energy 1 gram needs to absorb in
order to raise the temperature by 1 ˚C. (Note that the unit of “calorie” was defined by the amount of energy to
raise 1 gram of water by 1˚C-one calorie was defined by the specific heat capacity of water.
A substance with a very high heat capacity can absorb a lot of energy before the temperature rises. The sand in
summer time becomes hot much faster than the water in the ocean. Sand has a much lower heat capacity than
water. Also, the greater the mass of a substance is, the greater the energy that is needed to raise the
temperature of the entire substance. For example, it takes much longer to raise the temperature of a large pot
of water to the boiling point than it does for a small cup of water.
Using heat capacity in calculations:
If the specific heat capacity is the amount of energy needed to raise 1 gram by the 1˚C, then the equation that
relates heat capacity and energy absorbed by 1˚C, then the equation that relates heat capacity and energy
absorbed is ∆H = m x C x ∆T. The change in enthalpy is ∆H, and it uses energy units. The mass of the
substance is m, and it is measured in grams (g). The specific heat capacity is C, and its units are J/g˚C or
cal/g˚C. The change in temperature is ∆T, measured in ˚Celsius (˚C), and it is Tfinal – Tinitial.
SAMPLE PROBLEM: Read the following example carefully
Name:__________________________________________ Date:____________________ Period:_______
Critical Thinking Questions
5. When 435 J of heat is added to 3.4g of olive oil at 21˚C, the temperature increases to 85˚C. What is the
specific heat of the olive oil? (use the sample problem as a guide) Show your work
6. Explain why, when you take clothes that are still damp out of a clothes dryer, they don’t feel hot, but if
you wait until they’re dry to take them out, they feel quite hot. (hint: where is the energy from the
dryer going with the damp clothes? Where is it going with the dry clothes? Use the idea of specific heat
capacity in your explanation).
7. How many calories would be required to change the temperature of 750.0 g of water from 15˚C to
90˚C? Show your work
8. How many calories would be required to change the temperature of 250.0 g of aluminum from 15˚C to
75˚C? The specific heat of aluminum is 0.214 cal/g˚C. Show your work
9. Given 800.0 g of water at 22˚C, calculate the final temperature of the water after it absorbs 3600 calories.
10. Use specific heat capacity and energy to explain why farmers will spray their
orange trees with water to prevent frost damage during icy weather (i.e what
happen in terms of energy as water freezes).
11. The label on a box of apple pie warns that the “filling is hot”. When a freshly
baked apple pie comes out of the oven, both the filling and crust are at the same
temperature. However, the filling, which is mostly water can burn your tongue.
Why do you have to be careful not to burn your tongue?
Name:__________________________________________ Date:____________________ Period:_______
Example Questions for Practice
For a pure substance:




q = m x C x T
q = heat (joules or calories)
C = specific heat (J/(g° C) OR J/(cal° C))
m = mass (grams)
∆T = change in temp. (0C) = Tfinal – Tinitial
1 J=0.239 cal OR 4.184 J=1 cal
In-Class Example 1
 How much heat is required to raise the temperature of 250.0g of mercury 25°C? (Specific Heat
Hg = 0.14 J/(g• ° C) )
q=
C=
m=
∆T=
In-Class Example 2
 The temperature of a 95.4-g piece of copper increases from 25 ° C to 48.0 ° C when the
copper absorbs 849 J of heat. What is the specific heat of copper?
q=
C=
m=
∆T=
In-Class Example 3
 A swimming pool measuring 20.0cm x 12.5 cm is filled with water to a depth of 3.75cm. If the
initial temperature is 18.4°C, how much heat must be added to raise its temperature to
20.0°C. Density of water is 1.00g/mL…1 mL=1cm3. Specific heat 4.184 J/g·C°.
q=
C=
m=
∆T=
Name:__________________________________________ Date:____________________ Period:_______
H: PS#1 Relationship Between Heat Flow and
Temperature Change
1. Define thermochemistry
2. Define energy.
3.
Define heat and describe heat flow.
4.
Differentiate between the system and the surroundings.
5. What is an endothermic process? Describe the heat flow with regards to the system and surroundings.
6. What is an exothermic process? Describe the heat flow with regards to the system and surroundings.
7. What is the law of conservation of energy? What does that mean about the energy in the system and
surroundings?
8. If you want a hot pack to last for a very long time after it’s been heated, would you want it to have a
small specific heat capacity or a large specific heat capacity. EXLAIN your choice.
9. A diaper company has recently come out with training pants that let toddlers know they’re wet by
making them feel cold. Explain what happens when the pants get wet that causes them to make the
child feel cold.
10. What is the SI unit for energy? What are the conversion factors between that and a calorie?
11. Given the following equation, describe the meaning of each variable. Remember q = ∆H
Name:__________________________________________ Date:____________________ Period:_______
Show ALL work. Box final answer with units.
1. How many kiloJoules (KJ) of energy are needed to raise the temperature of 1.50 g of water from 20.0oC to
37.0oC? (Cwater=4.18 J/( g.C) )
2. The specific heats of three different liquid substances are listed as:
Carbon tetrachloride: 0.856 J/g.C
Benzene: 1.74 J/g.C
Acetic Acid: 2.05 J/g.C
An experimenter found that 1.42 kJ of heat energy raised the temperature of 19.70g of an unknown
liquid substance by 36.4oC. Based upon the substances listed above, what is the identity of the
unkown? (Show all work)
3. Mercury has a density of 13.546 g/cm3 and a specific heat of 0.139 J/g.C. How much energy in Joules is
released from 25.00 cm3 of Hg when it cools from the boiling point of Hg (357oC) to its freezing point
(-39oC)? scientific notation!
4. How many kJ are released from a 2.0 liter bottle of cola (2000 g) when it cools from 70oF (294K) to its
freezing point (273K)?
5. What (minimum) mass of glass (Cp= 0.749 J/g.C) at 26.0C is needed to absorb 5.00x104 Joules of heat
energy if its final temperature cannot exceed 275C?
6. What final temperature will 120.0 grams of benzene (Cp= 1.74 J/g.C) at 7.0C have after it absorbed 2.2kJ
of heat?
7. 3.0 kg of Osmium metal (Cp= 0.130 J/g.C) at 241 K is heated to 394 K. How much energy is needed for
this?
Name:__________________________________________ Date:____________________ Period:_______
8. 14.22 g of a substance absorbs 1.77 kJ of heat and undergoes a temperature change from –23.0C to
31.0C. What is the specific heat of the metal?
9. Calculate the amount of heat in kJ that was absorbed by a Sn (Cp= 0.220 J/g.C, D= 7.31 g/cm3) roof that
measures 32 feet by 20. feet if the sample is 0.0104 feet thick when the roof under goes a 15.0C
temperature change. Use scientific notation for final answer.
10. The density of gold is 19.3 g/cm3. What volume in cm3 of gold can absorb 2.3kJ of heat when undergoing a
5.0C T. It requires 0.128 J of heat to raise the temperature of 1g of Au 1C
Name:__________________________________________ Date:____________________ Period:_______
R: PS#1 Relationship Between Heat Flow and
Temperature Change
1. Define thermochemistry
2. Define energy.
3.
Define heat and describe heat flow.
4.
Differentiate between the system and the surroundings.
5. What is an endothermic process? Describe the heat flow with regards to the system and surroundings.
6. What is an exothermic process? Describe the heat flow with regards to the system and surroundings.
7. What is the law of conservation of energy? What does that mean about the energy in the system and
surroundings?
8. If you want a hot pack to last for a very long time after it’s been heated, would you want it to have a
small specific heat capacity or a large specific heat capacity. EXLAIN your choice.
9. A diaper company has recently come out with training pants that let toddlers know they’re wet by
making them feel cold. Explain what happens when the pants get wet that causes them to make the
child feel cold.
10. What is the SI unit for energy? What are the conversion factors between that and a calorie?
11. Given the following equation, describe the meaning of each variable. Remember q = ∆H
Name:__________________________________________ Date:____________________ Period:_______
Show ALL work. Box final answer with units.
1. How many kiloJoules (1000 J=1 KJ) of energy are needed to raise the temperature of 1.50 g of water
from 20.0oC to 37.0oC? (Cwater=4.18 J/( g.C) )
2. The specific heats of three different liquid substances are listed as:
3.
Carbon tetrachloride: 0.856 J/g.C
Benzene: 1.74 J/g.C
Acetic Acid: 2.05 J/g.C
An experimenter found that 1.42 kJ of heat energy raised the temperature of 19.70g of an unknown
liquid substance by 36.4oC. What substance that is listed above could this be? (Show all work)
4.
A sample of mercury has a mass of 338.65 g. The specific heat of mercury is 0.139 J/g.C. How much
energy in Joules is released the sample of Hg when it cools from the boiling point of Hg (357oC) to its
freezing point (-39oC)? scientific notation!
5.
How many kJ are released from a 2.0 liter bottle of cola (2000 g) when it cools from 21 C to its freezing
point 0C?
6.
What (minimum) mass of glass (Cp= 0.749 J/g.C) at 26.0C is needed to absorb 5.00x104 Joules of heat
energy if its final temperature cannot exceed 275C?
7.
What is the temperature change that 120.0 grams of benzene (Cp= 1.74 J/g.C) will undergo after it
absorbed 2.2kJ of heat?
8.
14.22 g of a substance absorbs 1.77 kJ of heat and undergoes a temperature change from –23.0C to
31.0C. What is the specific heat of the metal?
Name:__________________________________________ Date:____________________ Period:_______
POGIL Calorimetry
What is the relationship between heat energy and temperature?
Why?
When a substance is heated, the temperature of that substance increases. Will the same amount of energy
cause different substances to have identical temperature increases? Will the same amount of energy be needed
to cause identical temperature increases in different amounts of the same substance? This information is
essential to understanding the stability of chemical compounds, predicting equilibrium concentrations in
chemical reactions, and identifying conditions for a reaction to occur efficiently and safely. In this activity you
will learn how the energy change in a chemical reaction can be measured using a calorimeter and explore how
mass, temperature, heat energy, and the type of substance are related.
Model 1 – A Pot of Water
Critical Thinking Questions:
1. In Model 1, which container holds more grams of water?
2. Consider the process described in Model 1:
a. How many joules of energy were added to the saucepan?
b. How many joules of energy were added to the stockpot?
c.
In which container did the liquid gain more energy or did both gain the same amount?
Explain your reasoning.
d.
For each container, include whether the temperature is expected to increase, decrease or
remain the same after heating. Explain your reasoning.
Name:__________________________________________ Date:____________________ Period:_______
Model 2 – Experimental Data for Heating Water
Critical Thinking Questions:
3. In the data tables in Model 2, what does ∆T mean?
4. Which experiment in Model 2 illustrates how the amount of energy needed to achieve the same
temperature change depends on the mass of water?
5. Which experiment in Model 2 illustrates how different amounts of energy result in different
temperature changes when the mass of water is constant?
6. Why was it necessary to perform three experiments to find the relationships between mass,
temperature change, and energy?
Name:__________________________________________ Date:____________________ Period:_______
7. Refer to Experiment 1 in Model 2, and consider the relationship between the mass of water and the
observed temperature change when the same amount of energy is added.
a. Complete the statement below to show the relationship.
When the same amount of energy is added to water samples of different mass, the change in temperature gets
(smaller or larger) as the mass of the water increases.
b. Does the relationship stated in part a describe a direct or inverse relationship?
8. Refer to Experiment 2 in Model 2.
a. Write a grammatically correct sentence (like the one in Question 7a) to describe the relationship
between the observed temperature change and the energy required to heat water samples of identical
mass.
b. Does the relationship stated in part a describe a direct or inverse relationship?
9. In Experiment 2, should the value of the missing energy in the last row be larger or smaller than the
other energy values in that column?
10. Refer to Experiment 3 in Model 2.
a. Write a grammatically correct sentence to describe the relationship between the mass of water and the
energy required to produce the same temperature change in different water samples.
b. Does the relationship stated in part a describe a direct or inverse relationship?
11. Solve for the value of the constant “C” using data from Model 2 and the formula above. Your group
should calculate the value of C by using one set of data from each of the three experiments (include the
units of this constant). After you complete the three calculations, compare your value for the constant
with the value that other groups determined.
Show ALL calculations here:
Experiment #1
Experiment #2
Experiment #3
Name:__________________________________________ Date:____________________ Period:_______
12. The constant obtained in Question 11 is called the specific heat, (C). It is an intensive physical property
that has a different, characteristic value for every substance. What is the value for the specific heat of
water?
13. What do the units for specific heat mean? (Make sure your answer is a grammatically correct sentence.)
14. Using the equation above and the specific heat of water from Question 12, determine the values for the
missing data in the three experiments in Model 2. (Trial F in experiments 1–3.)
Show ALL calculation here:
Experiment #1
Experiment #2
Experiment #3
15. Use the equation above to calculate the following:
a. How much energy (q) is transferred when 30.0 g of water is cooled from 25.0 °C to 12.7 °C.
Show ALL calculation here:
b. Describe the significant difference between this value and the energy values shown in Model 2.
Information:
Name:__________________________________________ Date:____________________ Period:_______
Example #1: A 175 gram sample of a metal at 93.50C was added to 105 grams of water at 23.50C in a perfectly
insulated container. The final temperature of the water and metal was 33.80C. Calculate the specific heat of
the metal in J/g0C.
Heat lost by the metal
Mass of x specific heat x temp change
Metal
of metal
of metal
=
=
- heat gained by the water
- Mass of x specific heat x temp change
water
of water
of water
175 g
=
- 105 g x 4.184
g0C
=
- 105
=
0.43 J/g0C
x C x (33.8-93.5)0C
175 g x
C
x
59.70C
C = - (105)(4.184 J)(-10.3)
-(175 g)(59.70C)
= 0.4331 J/g0C
x
4.184 J
J
x (33.8-23.5)0C
x
10.3
Critical thinking questions: Whole-Class Practice
1. A 23.05 g of a metal alloy is heated to 1000C in boiling water. The alloy is placed in 50.00 g of water at
20.00C. The temp of the water increases to 29.50C. Determine the specific heat capacity of the alloy in
J/g0C. (∆t = tfinal - tinitial for each substance, the alloy and the water)
qalloy = malloy x ∆talloy x Calloy
= - qwater = mwater x ∆twater x Cwater
Use Example #1 and the whole-class practice from above as a guide to solve the following problems. Show all
work. Answers are given to you to see if you did the calculation correctly.
Critical Thinking Questions:
2. A student was given a sample of a silvery gray metal and told that it was either bismuth, specific heat
0.122 J/g0C, or cadmium, specific heat 0.232 J/g0C. The student measured out a 250 gram sample of the
metal, heated it to 96.00C and then added it to 98.5 grams of water at 21.00C in a perfect calorimeter.
The final temperature in the calorimeter was 30.30C. Use the student’s data to calculate the specific
heat of the metal sample and then identify the metal.
Name:__________________________________________ Date:____________________ Period:_______
3. A 245 gram sample of a metal at 99.50C was added to a 114 gram sample of water in a perfect
calorimeter. The original temperature in the calorimeter was 23.50C. The final temperature of the
metal-water mixture was 35.60C. Calculate the specific heat of this metal in joules per gram per Celsius
degree.
4.
175 grams of hot aluminum (100.°C) is dropped into an insulated
cup that contains 40.0 mL of ice cold water (0.0°C). Follow the
example above to determine the final temperature, x.
a. Set up an expression for the heat lost by the aluminum
(C=0.900 J/g·°C)
“shot” are these little pellets.
b. Set up an expression for the heat gained by the cold water.
c. Put the two expressions together (don’t forget to change one of the signs) and solve for x.
Name:__________________________________________ Date:____________________ Period:_______
Model 3—Lab: Which Metal Will Cook My Food Faster?
OBJECTIVE:
Determine the specific heat of several metal samples, and determine which would be the best metal to use for
cooking.
Pre-Lab Practice Calculations
1. How much heat is required to raise 57.1 grams of water from a temperature of 36.5 °C to
66.9 °C? (the specific heat of water = 4.184 J/g ºC)
2. What mass of magnesium is present if 250 J raises the temperature from 25 °C to 27 °C? (specific heat of
magnesium = 1.05 J/g ºC)
3. Determine the specific heat of a substance that absorbs 2500 joules of heat when a 50 gram sample
increases in temperature from 10.0ºC to 70.0ºC.
PART 1 - PROCEDURE: DETERMINING THE SPECIFIC HEAT OF A METAL
1. Take the mass of the metal obtained from your teacher. Be sure to record the letter on the metal.
2. Tilt a 250 mL beaker and gently place the metal inside of it.
3. Put enough water in the beaker to cover the metal, and place the beaker on a hot plate.
4. Allow the metal to heat up for 5-10 minutes. Continue to the next steps while you wait.
5. Setup a makeshift calorimeter by placing a coffee cup in a styrofoam cup.
6. Measure 50-75 mL of cool tap water in a graduated cylinder and pour this water into your calorimeter.
Record the exact amount (nearest 0.5 mL) as the mass of your water. (The density of water = 1.0 g/mL,
so the volume will equal the mass. Thus, 100 mL = 100 grams of water).
7. Measure the temperature (to the nearest 0.5ºC) of the water in the calorimeter. Record this as the
initial temperature of the water in your data table.
8. Check the thermometer of the hot water bath and record this temperature (to the nearest 0.5ºC) as the
initial temperature of the metal.
9. Using tongs, remove the metal from the hot water and immediately place it in the calorimeter.
Name:__________________________________________ Date:____________________ Period:_______
10. Monitor the temperature of the water in the calorimeter. Carefully use the thermometer to gently stir
the water. Record the highest temperature (to the nearest 0.5ºC) as the final temperature of the metal
and as the final temperature of the water in your data table.
11. Remove the metal from the cup, dry the metal, and repeat the procedure for another metal sample.
DATA:
Metal
Mass:
Symbol
m
Value
Trial 1
Value
Trial 2
Water
Symbol
Mass:
m
Initial
Temperature:
Initial
Temperature:
Final
Temperature:
Final
Temperature:
Change in
Temperature
(CALC)
Specific Heat:
(CALC)
Change in
Temperature:
(CALC)
Specific Heat:
Value
Trial 1
4.18 J/g C
Value Trial
2
4.18 J/g C
Calculations:
1) Using your data, determine the heat (q) gained by the water for both trials 1 and 2.
2) Using your answer from #1 as the heat (q) lost by the metal, determine the specific heat of your metal for
both trials 1 and 2.
Name:__________________________________________ Date:____________________ Period:_______
ANALYSIS QUESTIONS:
Please answer the following questions.
1. Using your calculated specific heat, determine the identity of your unknown metal. Check with your
teacher for the true identity. If specific heat is not enough to make a positive identification, you may wish to
calculate the density of your metal.
Metal
Aluminum
Copper
Iron
Lead
Magnesium
Nickel
Tin
Titanium
Zinc
Specific Heat (J/g ºC)
0.91
0.39
0.46
0.13
1.05
0.54
0.21
0.54
0.39
Density (g/mL)
2.70
8.92
7.87
11.34
1.74
8.91
7.31
4.51
7.14
2. Calculate the percent error for your specific heat for this experiment. Use your value as the experimental
and the value in the table as the accepted.
% error = absolute value (accepted value - experimental) x 100
accepted
3. So, which metal of your samples would be the best choice for your cookware? Assume you want to cook
your food the fastest. Do you have any pots or pans at home that contain this metal? Why or why not?
Explain your answer using the word specific heat..
4. We assumed that speed of cooking was our only concern. Explain how each of the following factors could
contribute to your decision of which pan would be the best for cooking.
a) Copper cookware is generally more expensive than Aluminum.
b) Aluminum is generally more chemically reactive than Copper
c) A cast iron pan is generally a lot thicker than an aluminum pan.
5. A beaker contains 125.0 grams of water at 22.0 °C. A small piece of metal with a mass of 38.2 grams is
hanging in another beaker of boiling water, measured at 102 °C. When the metal is removed and added to the
cooler water, the final temperature is 23.3 °C. What is the specific heat of the metal?
Name:__________________________________________ Date:____________________ Period:_______
Model 4—Thermochemcial Equations
Read pages 514-517 in your text to help answer the following questions.
Concept
Description/Explanation/Definition
(2-3 Sentence Minimum)
Thermochemical Equation
(in addition to the definition…write an example of both an
exothermic and endothermic thermochemical equation)
Heat of Reaction
(in addition to a definition, write the values of the heat of
reaction for an exothermic and endothermic reaction in the
textbook)
Figure 17.7a on page 515
Name:__________________________________________ Date:____________________ Period:_______
Concept
Description/Explanation/Definition
(2-3 Sentence Minimum)
Figure 17.7b on page 515
Heat of Combustion
Critical Thinking Question:
 Directions: Read sample problem 17.3 below and use it as a guide to solve question #7
5. The production of iron and carbon dioxide from iron(III) oxide and carbon monoxide is an
exothermic reaction. How many kilojouls of heat are produced when 3.40 mol of Fe2O3 reacts
with an excess of CO
Fe2O3 (s) + 3 CO (g)  2 Fe (s) + 3 CO2(g) + 26.3 KJ
Name:__________________________________________ Date:____________________ Period:_______
H PS#2: Calorimetry and Thermochemical Equations
1. A 15.0 gram sample of nickel metal is heated to 100.0 °C and dropped into 55.0 grams of water, initially at
23°C. Assuming that all the heat lost by the nickel is absorbed by the water, calculate the final temperature of
the nickel and water. (The specific heat of nickel is 0.444 J/g •°C)
2. 5.00 kg of a hot metal at 200.0C is added to 25.0 kg of water at 30.0C. What is the final temperature
of the metal? The specific heat of the metal is 0.800 J/g-C and for water it is 4.184 J/(gC). (36.2°C)
3. A small “coffee cup” calorimeter contains 110. g of water at 22.0C. A 100.g sample of lead is heated
to 90.0C and then placed in the water. The contents of the calorimeter come to a temperature of
23.9C. What is the specific heat of lead?
(0.132J/g .°C)
4. A 185 gram sample of copper at 98.00C was added to 102 grams of water at 20.00C in a perfectly
insulated calorimeter. The final temperature of the copper-water mixture was 31.20C. Calculate the
specific heat of copper using this data.
5. A chemistry student added 225 grams of aluminum at 85.00C to 115 grams of water at 23.00C in a
perfect calorimeter. The final temperature of the aluminum-water mixture was 41.40C. Use the
student’s data to calculate the specific heat of aluminum in joules/gram0C.
6. A 315 gram sample of tungsten at 92.50C was added to a 57.7 gram sample of water at 21.20C in a
perfect calorimeter. The final temperature of the tungsten-water mixture was 31.80C. Use this data to
calculate the specific heat of tungsten
Name:__________________________________________ Date:____________________ Period:_______
Thermochemical Equations
A thermochemical equation shows a balanced chemical equation and the corresponding enthalpy change, ΔH.
Indicate the correct answers for the following questions based on the thermochemical equation below.
Equation 1: Endothermic: 2C (s) + H2 (g)  C2H2 (g)
ΔH=226.6 kJ
OR
Equation 2: Exothermic:
2C (s) + H2 (g) + 226.6 kJ  C2H2 (g)
2CH4 (g) + O2 (g)  CO2 (g) + H2O (l)
ΔH=-870 kJ
OR
2CH4 (g) + O2 (g)  CO2 (g) + H2O (l) + 870 kJ
Look at the following example to help you complete question 7.
Example: How much heat is absorbed when 1 gram of ethylene gas, C2H2, is formed according equation 1
above?
1 g C2H2
1 mol C2H2
Molar mass of C2H2
= 0.038 mol C2H2
26.0 g C2H2
ΔH value from equation
0.038 mol C2H2 226.6 kJ
*Need moles because
you can obtain the
moles/kJ in the equation.
= 8.72 kJ
1 mol C2H2
Taken from the coefficient
7. How much heat would be absorbed when 4.2 moles of carbon react according to the first equation
above?
Name:__________________________________________ Date:____________________ Period:_______
8. Indicate the correct answers based on the following thermochemical equation:
2H2 (g) + O2 (g)  2H2O (l)
ΔH= -575 kJ
a. Is the reaction endothermic or exothermic? How do you know?
b. Which have the greater enthalpy: products or reactants?____________________________
c. Which are more stable, the products or reactants? ________________________________
c. How much heat is released when 1 mole of hydrogen gas reacts according to the thermochemcial
equation?
e. How many grams of liquid water are formed when 159.7 kJ of heat is released?
11. The standard enthalpy of formation of sodium, chloride is -411 kJ/mol. Write the thermochemical equation
for the formation of sodium chloride from sodium and chlorine under standard conditions.
12. Calculate the energy required to produce 5.0 moles of Ca(OH)2
CaO (s) + H2O (l)  Ca(OH)2
ΔH= -65.2 kJ
Name:__________________________________________ Date:____________________ Period:_______
R PS#2: Calorimetry and Thermochemical Equations
1. A small “coffee cup” calorimeter contains 110. g of water at 22.0C. A 100.g sample of lead is heated
to 90.0C and then placed in the water. The contents of the calorimeter come to a temperature of
23.9C. What is the specific heat of lead?
(0.132J/g .°C)
2. A 185 gram sample of copper at 98.00C was added to 102 grams of water at 20.00C in a perfectly
insulated calorimeter. The final temperature of the copper-water mixture was 31.20C. Calculate the
specific heat of copper using this data.
3. A chemistry student added 225 grams of aluminum at 85.00C to 115 grams of water at 23.00C in a
perfect calorimeter. The final temperature of the aluminum-water mixture was 41.40C. Use the
student’s data to calculate the specific heat of aluminum in joules/gram0C.
4. A 315 gram sample of tungsten at 92.50C was added to a 57.7 gram sample of water at 21.20C in a
perfect calorimeter. The final temperature of the tungsten-water mixture was 31.80C. Use this data to
calculate the specific heat of tungsten
5. A 15.0 gram sample of nickel metal is heated to 100.0 °C and dropped into 55.0 grams of water, initially at
23°C. Assuming that all the heat lost by the nickel is absorbed by the water, calculate the final temperature of
the nickel and water. (The specific heat of nickel is 0.444 J/g •°C)
Name:__________________________________________ Date:____________________ Period:_______
Thermochemical Equations
A thermochemical equation shows a balanced chemical equation and the corresponding enthalpy change, ΔH.
Indicate the correct answers for the following questions based on the thermochemical equation below.
Equation 1: Endothermic: 2C (s) + H2 (g)  C2H2 (g)
ΔH=226.6 kJ
2C (s) + H2 (g) + 226.6 kJ  C2H2 (g)
OR
Equation 2: Exothermic:
2CH4 (g) + O2 (g)  CO2 (g) + H2O (l)
ΔH=-870 kJ
OR
2CH4 (g) + O2 (g)  CO2 (g) + H2O (l) + 870 kJ
Look at the following example to help you complete question 6.
Example: How much heat is absorbed when 1 gram of ethylene gas, C2H2, is formed according equation 1
above?
1 g C2H2
1 mol C2H2
Molar mass of C2H2
= 0.038 mol C2H2
26.0 g C2H2
ΔH value from equation
0.038 mol C2H2 226.6 kJ
*Need moles because
you can obtain the
moles/kJ in the equation.
= 8.72 kJ
1 mol C2H2
Taken from the coefficient
6. How much heat would be absorbed when 4.2 moles of carbon react according to the first equation
above?
Name:__________________________________________ Date:____________________ Period:_______
7. Indicate the correct answers based on the following thermochemical equation:
2H2 (g) + O2 (g)  2H2O (l)
ΔH= -575 kJ
a. Is the reaction endothermic or exothermic? How do you know?
b. Which have the greater enthalpy: products or reactants?___________________________________
c. Which are more stable, the products or reactants? ________________________________________
d. How much heat is released when 1 mole of hydrogen gas reacts according to the thermochemcial
equation?
e. How many grams of liquid water are formed when 159.7 kJ of heat is released?
11. The standard enthalpy of formation of sodium, chloride is -411 kJ/mol. Write the thermochemical equation
for the formation of sodium chloride from sodium and chlorine under standard conditions.
12. Calculate the energy required to produce 5.0 moles of Ca(OH)2
CaO (s) + H2O (l)  Ca(OH)2
ΔH= -65.2 kJ
Name:__________________________________________ Date:____________________ Period:_______
Lab: Making Hot and Cold Packs
Information:
Hot and cold packs are frequently used by athletes and trainers to treat minor injuries, such as,
inflammations, sprains, muscle spasms, head-aches, etc. Hot and cold packs consist of two separate
compartments, one containing water and the other containing a salt. When you break the seal between the
chambers and shake the pack vigorously, the two compartments combine and the salt dissolves in the water.
This generates a chemical reaction. Depending on the type of salt, this reaction can either release or absorb
heat energy. When a reaction occurs that releases heat, it is referred to as an exothermic reaction (the prefix
exo is Latin for “out of”). In contrast, if the reaction requires heat to occur, it is referred to as an endothermic
reaction (the prefix endo is Latin for “into”).
A hot pack is produced if an exothermic reaction occurs as the salt and water mix and heat energy is
released because this process will raise the temperature of the contents in the pack. A cold pack is produced if
an endothermic reaction occurs as the salt and water mix and heat energy is absorbed because this process
will lower the temperature of the contents in the pack. The amount of heat that is released or absorbed by the
packs depends on the concentration of water and salts.
To fully understand how a hot or cold pack works, we
must examine how the particles interact when a salt dissolves
Figure 1: Solvation of a Salt
in water (see figure 1). The salt crystal is made of positive and
negative ions that hold it together (eg. NaCl or Na+ and Cl-)
by electrostatic ion-ion attractions (the attraction of opposite
charges). When a salt dissolves in water, the electrostatic
attractions (attractions between positive and negative
charges) between the ions are broken and each ion forms new
electrostatic interactions with the water molecules (see figure
1). Whether this process creates an endothermic or
exothermic reaction depends on the balance between the ionion forces of the solid salt that must be overcome and the
dipole-ion forces that stabilize the dissolved ions in solution.
Purpose: To measure the solubility as different compounds dissolve in water and determine the specific heat
of the solutions.
Materials:
goggles
distilled water
thermometer
stirring rod
CaCl2
graduated cylinder
foam cup as a calorimeter
balance
NH4Cl
Name:__________________________________________ Date:____________________ Period:_______
Procedure:
1.
Put on goggles. Measure 100.0 mL of distilled water at room temperature and pour it into
the plastic foam cup. Record the mass of the water by recalling that the density of water is
1g/mL. Record the temperature of the water in the data table to the nearest 0.1C. Do not
remove thermometer from the cup, but be careful that it does not tip over.
2.
Using the electronic balance, place a piece of paper on the balance and zero the balance.
Measure out 8-10 grams of ammonium chloride on the piece of paper. Record the mass to
the nearest 0.01 g in the table below (*if the balance permits).
3.
Without removing the thermometer from the cup, shake the NH4Cl from the paper into the
water and stir gently with the stirring rod until the solid is dissolved. CAUTION: Both of
the solutions in the lab are irritating to the skin. Avoid contact with them.
4.
Make sure that the bulb of the thermometer is fully immersed in the liquid. If the
temperature rises, record the highest temperature reached. If the temperature falls, record
the lowest temperature reached.
5.
Dispose of the solution by pouring it down the drain, followed by plenty of water. Rinse the
cup, dry and return it and the thermometer to the lab bench.
6.
Repeat steps 2-5 using calcium chloride.
Observations:
Data Table
Solute
Solute
Mass (g)
Mass of
Water (g)
Mass of
Solution
(g)
Initial T
(C)
Final T
(C)
(+/-) T
(C)
Exothermic
or
Endothermic
NH4Cl
CaCl2
Calculations for NH4Cl
1. Calculate the change in temperature. T = Tf ─ Ti
2. Calculate the heat absorbed or released by the solution. The specific heat of water is 4.184J/gC.
qsur = (mass of the solution) x (T of the water) x (specific heat of the water)
3. How much heat was released/absorbed (circle one) from/into the reaction? (qrxn = -qsur)
qrxn =
Name:__________________________________________ Date:____________________ Period:_______
4. Using the equation from Model 3, calculate the specific heat of NH4Cl.
5. Using the periodic table, calculate the molar mass of the solute.
6. How many moles of the solute were used in the reaction?
7. Calculate the molar heat of solution (H) in (+/-) kJ for the solute from the formula:
Heat of the reaction
q
∆H = Molar heat of solution =
= rxn
moles solute dissolved
mole
8. Calculate the percent error of your experimental value.
% error =
accepted value - experimental value
x 100% =
accepted value
The accepted value for H for NH4Cl is +14.6 kJ/mol.
The accepted value for H for CaCl2 is –82.8 kJ/mol.
Calculations for CaCl2
1. Calculate the change in temperature. T = Tf ─ Ti
2. Calculate the heat absorbed or released by the solution. The specific heat of water is 4.184J/gC.
qsur = (mass of the solution) x (T of the water) x (specific heat of the water)
3. How much heat was released/absorbed (circle one) from/into the reaction? (qrxn = -qsur)
qrxn =
Name:__________________________________________ Date:____________________ Period:_______
4. Using the equation from Model 3, calculate the specific heat of CaCl2.
5. Using the periodic table, calculate the molar mass of the solute.
6. How many moles of the solute were used in the reaction?
7. Calculate the molar heat of solution (H) in (+/-) kJ for the solute from the formula:
Heat of the reaction
q
∆H = Molar heat of solution =
= rxn
moles solute dissolved
mole
8. Calculate the percent error of your experimental value.
% error =
accepted value - experimental value
x 100% =
accepted value
The accepted value for H for NH4Cl is +14.8 kJ/mol.
The accepted value for H for CaCl2 is –81.3 kJ/mol.
Critical Thinking: Analysis and Conclusions
1. When sodium chloride dissolves in water, the ions dissociate. The equation for this reaction is
NaCl(s)  Na+(aq) + Cl-(aq) Write similar ionic equations to show the dissociation in water of each of
the solutes used in the investigation.
Rxn with NH4Cl:______________________________________________________________________
Rxn with CaCl2:______________________________________________________________________
2. Which reaction was:
a. Exothermic: _______________________
b. Endothermic_______________________
Name:__________________________________________ Date:____________________ Period:_______
3. Rewrite each of the ionic equations from Question 1 showing the molar heat of solution as a reactant
or a product. For example: CH4 + O2  CO2 +H2O + ΔH
_______________________________________________________________________________
_______________________________________________________________________________
4. When the reactants get colder in an endothermic reaction, what has happened to the heat energy?
_____________________________________________________________________________
5. Is the change in enthalpy positive or negative for an exothermic reaction? Explain. __________
_____________________________________________________________________________
6. Suggest two uses for these solution reactions in sports injuries or camping.
_______________________________________________________________________________
_______________________________________________________________________________
7. Which solids from this investigation could be used in each of your answers in the previous question?
Explain your reasoning.
_______________________________________________________________________________
_______________________________________________________________________________
8. How could you reduce experimental errors in this investigation? Explain your reasoning.
_______________________________________________________________________________
_______________________________________________________________________________
Name:__________________________________________ Date:____________________ Period:_______
Lab
Why do some salts produce hot packs and some produce cold packs? In the case of an ionic solid dissolving in
water, the overall energy change is the net result of three processes—
1) The energy required to break the attractive forces between ions in the crystal lattice (ΔH1 = + C kJ/mole)
2) The energy required to disrupt intermolecular forces between water molecules (ΔH2 = + D kJ/mole)
3) The energy released when the dissociated (free) ions form ion-dipole attractive forces with the water
molecules (ΔH3 = − F kJ/mole).
The overall process can be represented by the following equation.
ΔHsoln = ΔH1 + ΔH2 + ΔH3 = (+ C + D − F) kJ/mole
1) If the amount of energy released in the formation of hydrated ions (ΔH3) is greater than the amount of energy
required to separate the solute and solvent particles (ΔH1 + ΔH2), then the sum (ΔHsoln) of the energy changes
will be negative and the solution process exothermic (releases heat).
2) If the amount of energy released in the formation of hydrated ions is less than the amount of energy
required to separate the solute and solvent particles, then the sum of the energy changes will be positive and
the solution process endothermic (absorbs heat).
Guided-Inquiry Design and Procedure
1. Review the calorimetry procedure and answer the following questions:
a. What data is needed to calculate the enthalpy change for a reaction?
b. Identify the variables that will influence the experimental data.
c. What variables should be controlled (kept constant) during the procedure?
d. The independent variable in an experiment is the variable that is changed by the experimenter,
while the dependent variable responds to or depends on the changes in the independent
variable. Name the independent and dependent variables in a calorimetry experiment to
determine the molar heat of solution.
e. Discuss the factors that will affect the precision of the experimental results.
2. Form a working group with four students per group. One pair of students in the group should study
the three solids in Set A, while the other pair studies Set B.
3. Working collaboratively with the general procedure provided in the Introductory Activity, design and
carry out experiments to determine the heat of solution for each solid. Be sure to review all safety
precautions with your instructor before starting.
4. Extrapolating from the information collected, predict which solid(s) could be used in an effective hand
warmer meeting the following requirements:
Name:__________________________________________ Date:____________________ Period:_______
• The hand warmer must contain 10 g of an ionic solid and an inner pouch filled with 40 mL of water.
• Activating the hand warmer must increase the temperature of the resulting solution by at least 20 °C.
5. Review the cost information shown below and consult the MSDS for each potential hand warmer.
Propose the optimum design for the most cost-effective hand warmer that is nontoxic and least harmful
to the environment.
6. With your instructor’s permission, verify the design and demonstrate the use of your hand warmer.
Name:__________________________________________ Date:____________________ Period:_______
Lab –
General
Name:__________________________________________ Date:____________________ Period:_______
Name:__________________________________________ Date:____________________ Period:_______
Name:__________________________________________ Date:____________________ Period:_______
Name:__________________________________________ Date:____________________ Period:_______
Name:__________________________________________ Date:____________________ Period:_______
Name:__________________________________________ Date:____________________ Period:_______