The Hand Warmer Challenge
Introduction:
Have your fingers ever been so cold they felt numb? Wouldn’t it be great if you could
generate heat to warm your hands up anytime you want to? That’s exactly what a “hand
warme
...
The Hand Warmer Challenge
Introduction:
Have your fingers ever been so cold they felt numb? Wouldn’t it be great if you could
generate heat to warm your hands up anytime you want to? That’s exactly what a “hand
warmer” does. Hand warmers are small packets that people put inside gloves or mittens on
cold days to keep their fingers warm. They are very popular with people who work outside in
winter or engage in winter sports. One type of hand warmer contains water in one section of
the packet and a soluble substance in another section. When the packet is squeezed the
water and the soluble substance are mixed, the solid dissolves and the packet becomes
warm. In this experiment, you will learn how a hand warmer works and use chemistry to
design an effective, safe, environmentally benign (no harm is done to the environment
through their use), and inexpensive hand warmer.
In this lab, you are challenged to use chemistry to design an effective, safe, environmentally
benign, and inexpensive hand warmer. The ideal hand warmer increases in temperature by
20 oC (but no more) as quickly as possible, has a volume of about 50 mL, costs as little as
possible to make, and uses chemicals that are as safe and environmentally friendly as
possible. You will carry out an experiment to determine which substances, in what amounts,
to use in order to make a hand warmer that meets these criteria.
Breaking bonds and particulate attractions absorb energy from the surroundings, while
forming new bonds and particulate attractions release energy to the surroundings. When an
ionic solid dissolves in water, ionic bonds between cations and anions in the ionic solid and
hydrogen bonds between water molecules are broken, and new attractions between water
molecules and anions and water molecules and cations are formed. The amount of energy
required to break these bonds and form new ones depends on the chemical properties of the
particular anions and cations. Therefore, when some ionic solids dissolve, more energy is
required to break the cation–anion bonds than is released in forming the new water–ion
attractions, and the overall process absorbs energy in the form of heat. When other ionic
compounds dissolve, the converse is true, and the bond making releases more energy than
the bond breaking absorbs, and therefore the process overall releases heat. When heat is
absorbed, the enthalpy change, q, is endothermic, and the enthalpy change is positive.
When heat is released, the change is exothermic, and the value of q is negative. Recall that
heat (q) can be calculated by:
q = mCsΔT
The entropy (disorder) change of solution formation is always positive, regardless of whether
it is endothermic or exothermic, because solutions are much more disordered than are the
pure solute and solvent from which they are made. This positive entropy change is
thermodynamically favorable.
In this experiment, you will collect data that will allow them to calculate the change of
enthalpy of dissolution (also called the “heat of solution,” with symbol ΔHsoln, and units of
kJ/mol solute) occurring in aqueous solution. The data necessary to calculate the heat of
solution can be obtained using a device called a calorimeter. A calorimeter is a container
used to determine the enthalpy change that occurs during a process. Calorimetry is an
important technique in chemistry, and chemists often work with devices called bomb
calorimeters. For home or classroom experiments, however, a coffee cup calorimeter is
Your Objective for this Investigation: Students will design and execute an experimental
procedure to determine which of three ionic compounds is most suitable for use in a hand
warmer.
Materials Available:
Materials Quantity
Thermometers 1
Balance 1
Graduated Cylinder, 100 mL 2
Coffee Cup Calorimeters 1
Hot Plate 1
Sodium Acetate, NaC2H3O2 5 – 10 g
Magnesium Sulfate, MgSO4 5 – 10 g
Calcium Chloride, CaCl2 5 – 10 g
Lithium Chloride, LiCl 5 – 10 g
Ammonium Nitrate, NH4NO3 5 – 10 g
Sodium Chloride, NaCl 5 – 10 g
Pre-Lab:
1. What are endothermic and exothermic reactions?
2. Go explore the following animation. Take notes on the topics that are addressed here:
http://group.chem.iastate.edu/Greenbowe/sections/projectfolder/flashfiles/thermochem/solutionSalt.html
Describe the changes you observe in the animation, including changes in the bonds and
particulate attractions and changes in the amount of disorder in the system.
3. When sodium chloride is dissolved in water, the temperature of the resulting solution is lower than
the temperature of the water before the salt dissolves. How can this result be explained based on the
bond breaking and bond making that is occurring?
4. Why do some salts, such as sodium chloride, dissolve spontaneously even though the process is
endothermic overall?
5. When some ionic salts are dissolved in water, the temperature of the resulting solution is higher
than the temperature of the water before the salt dissolves. What do you think determines whether
the resulting solution is cooler or warmer than the starting water?
6. What is a coffee cup calorimeter? Research how these work and make a drawing of how
calorimeters work.
7. How do you use a coffee cup calorimeter to calculate heat?
Procedure:
Part I: Calorimetry Practice
1. Assemble your calorimeter as you drew in question 7 of the pre-lab. Measure out exactly
100.0 mL water in a graduated cylinder and pour into your calorimeter. Measure and record
the temperature of the water.
2. Measure 5.00 g magnesium sulfate anhydrate solid into a plastic cup. While monitoring
the temperature of
the water, quickly add all of the magnesium sulfate to the calorimeter.
3. Record the highest temperature reached. Dilute the resulting solution with water and
dispose of it per your teacher’s instructions.
4. Repeat the experiment. For each trial find the temperature change of the water per gram
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https://www.coursehero.com/file/14453268/Lab-9-The-Hand-Warmer-Challenge-1-/of magnesium sulfate, and average this value for the two trials. Dispose of materials upon
completion of this step.
Part II: Calorimeter Calibration Practice
5. Place a 100.0 mL sample of water in a clean, dry 150 mL beaker. Heat with occasional
stirring to approximately 50 oC. Remove the beaker from the hot plate and place on the lab
bench.
6. Meanwhile, place exactly 100.0 mL of cool water (approximately 20 oC) in the clean, dry
calorimeter.
7. Measure the temperature of the hot water and the cold water and record, then
immediately pour the entire hot water sample into the calorimeter and quickly put on the
cover. Wait 15 seconds then take a temperature reading.
8. Repeat this determination twice.
Part III: Designing and Implementing Lab
9. Rank the solids you are given from least to most expensive.
Substan
ce
2012 Cost per 500
g ($)
NaCl 3.95
CaCl2 6.55
NaC2H3O2 12.90
Na2CO3 6.15
LiCl 32.75
NH4NO3 9.05
10. Based on the pre-lab work and the information provided, formulate a research question
that you wish to study in regards to hand warmers and salts:
11. Based on the pre-lab work and the information provided, formulate a hypothesis to your
research question:
12. Using the research from the pre – lab, work with your group to design a procedure to compare
the solids in terms of the heat released or absorbed when they dissolve and include what materials
and equipment you will use. You must include the safety precautions you will take.
IMPORTANT PROCEDURE TIPS:
(i) Be sure to keep detailed records of the amounts of substances used and the starting and
ending temperature as you will need it later to determine the amount of solid to use in your
hand warmer.
(ii) You will receive a maximum of 10 g of each solid for this part
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