The Hand Warmer Design Challenge: Where Does the Heat …



The Hand Warmer Design Challenge: Where Does the Heat Come From?

The goal of this lab is to design a safe, effective, environmentally benign, and inexpensive hand warmer that will increase the temperature of water by 20 °C (but no more) as quickly as possible with a volume of about 50 mL.

Introduction/Background

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 any time 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, students will learn how a hand warmer works and use chemistry to design an effective, safe, environmentally benign, and inexpensive hand warmer.

Breaking bonds and particulate attractions absorbs energy from the surroundings, while forming new bonds and particulate attractions releases 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 cations and anions. 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 energythan the bond-breaking absorbs, and therefore the process overall releases heat.

Day One:

Prelab activities and questions

1) View the Heat Transfer between Metal and Water tutorial at: flashfiles/thermochem/heat_metal.html (link at )

Simulation Part 1

(a) Choose silver

(b) Select the following variables: mmetal = 50.0g, T1metal = 100.0oC, mwater = 100.0g, T1water = 25.0oC.

(c) Run the simulation (make sure you record all data, including the highest temperature reached) and show the calculations to determine the qwater and qmetal.

Simulation Part 2

b) Choose Metal X.

c) Perform an experiment to determine the specific heat capacity of the metal.

d) Record data and calculate the specific heat capacity of the metal.

e) Repeat your experiment using metal X again, but select a different set of variables.

f) Record data and calculate the specific heat capacity of the metal.

g) How do the specific heat capacities from c and e compare? Is this the expected result?

h) What are three possibilities for the identity of your metal? (Need a specific heat capacity list? Find one at ; link at )

9) View the Energy Exchanges Associated with Dissolving Salts in the Water tutorial at: flashfiles/thermochem/heat_soln.html (link at )

(Show work for zinc sulfate in left hand side, ammonium chloride in right hand side)

j) Choose 1.00 g of zinc sulfate and 20.00 mL of water.

k) Perform the experiment.

l) Record the masses and temperatures.

m) Calculate the change in energy that occurred in joules per gram of chemical used.

n) Repeat this process, but double the amount of zinc sulfate.

o) How did the calculation change?

p) Repeat this process but use 1.00 g of zinc sulfate and 40.00 mL of water.

q) How did the calculation change?

r) Is zinc sulfate dissolving in water an exothermic or endothermic change? Explain.

s) Repeat steps a-i with ammonium chloride.

20) View the Dissolving of Salt animation at: flashfiles/thermochem/solutionSalt.html (link at )

u) This animation shows the ionic compound sodium chloride dissolving in water. What do the green blobs represent? What do the silver blobs represent?

v) Describe the changes you observe in the animation, including changes in the bonds and particulate attractions.

w) When some ionic compounds are dissolved in water, the temperature of the resulting solution is higher than the temperature of the water before the salt dissolves, while with other ionic compounds, the resulting solution has a lower temperature than before the dissolving. What do you think determines whether the resulting solution is cooler or warmer than the starting water?

24) What is the formula for calculating energy changes using specific heat? Identify what each variable in the formula stands for and its units.

25) What does endothermic mean? If dissolving a certain salt is an endothermic process, will the final temperature be higher or lower than the start?

26) What does exothermic mean? If dissolving a certain salt is an exothermic process, will the final temperature be higher or lower than the start?

27) What is the sign (+ or -) for an endothermic change? Why is that the sign?

28) What is the sign (+ or -) for an exothermic change? Why is that the sign?

29) Look at the following chart of prices for the solids we will use in this experiment:

|Substance |Formula |Safety |2012 Cost per 500 g |

|sodium chloride | | |$3.95 |

|calcium chloride | | |$6.55 |

|sodium acetate | | |$12.90 |

|sodium carbonate | | |$6.15 |

|lithium chloride | | |$32.75 |

|ammonium nitrate | | |$9.05 |

a) Record the formula for each solid.

b) Rank the solids from most expensive to least expensive

(c) Search for a MSDS or SDS sheet for each and compare. Assign a safety ranking with 1 the most safe and 6 the least safe.

Day 2 Lab Experiment

  Explanation to strengthen student understanding:

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. The entropy 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.

Practice with instrumentation and procedure

In this experiment, students will collect data that will allow them to calculate the change of enthalpy of dissolution (also called the “heat of solution,” with symbol ÄHsoln, 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 sufficient to make rough measurements. This exercise will give students practice assembling and using a calorimeter so that they can use one to help them determine which solid is best to use in a hand warmer. It will also allow students to calibrate their calorimeters with a process that supplies a known amount of heat. This calibration process allows students to determine the amount of heat the calorimeter itself absorbs as the temperature of the materials inside it change, a value known as a calorimeter constant.

Calorimetry Training Activity and Calorimeter Calibration

Discussion: Calorimetry is a process that joins the ideas of thermochemistry and stoichiometry. By tracking the amount of energy given off or taken in during a certain instance of a reaction, the more formal value, Change in Enthalpy (DH) can be determined.

Heat Capacity of the Calorimeter:

Calibration of calorimeter: (assume the density of water is ~1.00 g/mL for all temps)

1. Measure 100.0 mL of cold water and pour the cold water into the calorimeter as shown in class.

2. Heat 100.0 mL of how water to about 50.00oC

3. Measure the temperature of the cold water

4. Measure the temperature of the hot water

5. Pour the hot water into the calorimeter

6. Stir and record the equilibrium temperature in the calorimeter

|Mass Cold Water |Temp Cold Water |Mass Hot Water |Temp Hot Water |Final Temp Mix |

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|Calorimeter Constant Calculations Table |

|Temp. change hot water, ?Thot(°C) | |

|Tmix – Ti hot water | |

|Enthalpy change hot water, qhot (J) | |

|Q = mcDT | |

|Temp. change cold water, ?Tcold (°C) | |

|Tmix – Ti cold water | |

|Enthalpy change hot water, qcold (J) | |

|Temp. change calorimeter, ?Tcal (°C) | |

|Tmix – Ti cold water | |

|Enthalpy change calorimeter, qcall (J) | |

|¦¦qhot¦- qcold ¦ | |

|Calorimeter constant, C, (J/°C) | |

|Qcal / DTcal | |

Heat of Solution of Magnesium Sulfate: MgSO4(s) + H2O(l) ( Mg2+(aq) + SO42-(aq)

Materials:

Styrofoam cup Thermometer Balance Graduated cylinder

Procedure:

1. Measure out 25 ml of water and put it into the Styrofoam cup.

2. Determine the temperature of the water and record it on the data table.

3. Mass out approximately 2 grams of magnesium sulfate. Record the exact mass on the data table.

4. Place the magnesium sulfate into the water and swirl the liquid.

5. Monitor the temperature of the solution over a period of two minutes. Record, on the data table, the coldest temperature measured.

6. Dump the solution down the drain, rinse and dry the cup.

Data:

|Initial temp of the water |°C |

|Mass of the MgSO4 |g |

|Coldest temp of the solution |°C |

Calculations:

1. Determine the q for the water: qwater = mCDT + CcalDT

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1. What is the value of qsoln? qsoln = -qwater

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1. Calculate the number of moles of magnesium sulfate used.

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1. How many kilojoules of energy were transferred? (J(kJ)

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5. What is the DHsoln? Qsoln / mol

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Group Experiment

Procedure - Day 2 (List in your lab book):

You have been assigned the chemical: _____________________________________

Determine the DHsoln of your assigned solid and post your results to the group data table. Summarize your procedure and calculations for this here:

|Group Data Table |

|Solid |Group 1 DHsoln |Group 2 DHsoln |Average DHsoln |

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Post Lab Questions:

Looking at the class’s data, considering costs, and considering safety concerns, which of the chemicals would make the best hand warmer? Explain.

Ignoring safety concerns, which chemical gives the most temperature change for the money? Explain.

Were the dissolving process for each exothermic, endothermic, or neither? Explain.

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