5.1: Calorimetry/Thermochemistry Lab Procedure

Learning Objectives

• Determine whether a reaction is endothermic or exothermic.
• Determine the best ionic compound of to use in a heat pack for treating frostbite based on your experimental results.
• Calculate the average heat capacity of your calorimeter.
• Calculate the specific heat of a metal.

 

Background

The human body works best within a very narrow temperature range. A temperature drop of as little as 2o C in the body’s core causes classic hypothermia symptoms such as mental difficulties and loss of physical coordination. Much more extreme temperature drops in the extremities may be survived, but can lead to frostbite if the flesh freezes. Victims of hypothermia require immediate treatment, and in outdoor situations the treatment is often warmth provided by portable heat sources, such as heat packs.

Heat packs are available that produce warmth through various chemical reactions. Such heat packs are convenient because they only release heat when triggered. One common heat pack contains an internal pouch of water and a solid powder. Once the pouch of water is broken open, there is an exothermic reaction between the water and the powder. These heat packs have limitations. For example, they do not work well in extreme cold because the water within the pack will freeze.

In the coldest environments, heat packs are available that contain only the powder in a resealable waterproof sack. When heat is required, the sack can be opened and any aqueous solution poured inside. The sack is resealed and the reaction produces heat.

 

Reaction Enthalpy

An important part of chemistry involves studying energy changes that occur during chemical reactions. These energy changes are of fundamental importance in understanding the “driving force” of a chemical reaction. The most common way energy is exchanged between a chemical system and the environment is by evolution or absorption of heat (q). The change in heat energy accompanying a chemical reaction is known as enthalpy change, ΔH. By convention, reactions in which heat is absorbed are labeled endothermic and have positive values of ΔH; reactions in which heat is released are labeled exothermic and have negative values of ΔH. Cold packs used in athletics are familiar to many sports enthusiasts. In order to derive coldness from the pack, a plastic packet of water is broken inside another packet containing a solid salt such as NH4NO3. In this case, the enthalpy of solution, i.e., heat absorbed when a substance dissolves, is endothermic indicating that heat is absorbed as the salt dissolves. Thus, the enthalpy of solution is designated with a positive sign since energy is absorbed or added.

\[\ce{NH4NO3(s) -> NH4NO3(aq)} \; \Delta H = + 25.7 kJ \label{1}\]

On the other hand, gas stoves produce heat through the combustion (burning in oxygen) of methane, (see Equation 1. Since heat is produced by the combustion, the reaction is exothermic and enthalpy of combustion (ΔH), i.e., heat released during combustion, must be negative. In fact, ΔH for the methane combustion in is –890.4 kJ (see below). It is important to realize that ΔH is related to the coefficients in the balanced equation. Thus, –890.4 kJ of heat is released per every 1 mole of CH4 that reacts but per every two moles of H2O that is formed.

\[\ce{CH4(g) + O2(g) -> CO2(g) + 2H2O(g)} \; \Delta H = -890.4 kJ \label{2}\]

The negative sign serves to reinforce the fact that heat is produced rather than absorbed or that this is an exothermic reaction. We can use stoichiometry to calculate the enthalpy of any amount of a reactant.

 

Heat and Temperature

If an object (such as a pot of water) is positioned to absorb the heat given off during a combustion reaction, then the temperature of the object will change as follows

\[q = m \times c \times \Delta T \label{MCAT} \]

Where q = the amount of heat absorbed by the object in Joules

m = the mass of the object being heated in grams

c = the specific heat of the object being heated

ΔT = the change in temperature of the object = final temperature minus the initial temperature = Tf – Ti

The specific heat is different for different substances. The specific heat or heat capacity is low for objects which are easily heated or cooled with minimal energy input. Metals usually have low specific heats because their temperature changes very quickly. Water has a high specific heat because it takes a lot of energy to change the temperature of a sample of water. For example:

Example specific heat values

Substance	Specific Heat (J g^-1C^-1)
Water	4.18
Air	1.01
Aluminum	0.897
Granite	0.790

 

Calorimetry

 

Figure \(\PageIndex{1}\): A simple illustration of a typical calorimeter.

Many experiments in thermochemistry involve a calorimeter. A calorimeter, like the one above, is simply a container that insulates a reaction from the surrounding environment so change in temperature as the reaction proceeds place can be measured accurately measured as independent of the environmental temperature. In lab, you will use two Styrofoam cups as your calorimeter. Ideally any calorimeter would be able to completely maintain the system without losing heat to the surroundings. Unfortunately, no calorimeter is perfect and heat is always lost to the surroundings. It is possible to determine how much heat is lost by the calorimeter by mixing hot water and room temperature water. The heat lost by the hot water is equal to the heat gained by the room temperature water.

\[\left | -q_{hot} \right | = \left | q_{ambient} \right | + \left | q_{calorimeter} \right |  \label{QQQ} \]

In order to calculate the heat lost to the calorimeter you will calculate q for both the hot water and the room temperature water (using \ref{MCAT}) and then subtract to find the difference. It is also possible to view how heat can be transferred from one object to another. In a perfect system (where no heat is lost to the calorimeter) the heat lost by the hot object would be gained by the cold object such as:

\[q_{lost} = -q_{gained} \label{qlost}\]

which could be rearranged to give:

\[q_{hot\; water} = -q_{cold\; water} \label{qhot}\]

as well as replacing q with \ref{QQQ} to give

\[mc(T_f - T_i)_{hot\; object} = -mc(T_f - T_i)_{cold\; object} \label{complete}\]

Although usually heat is lost to the calorimeter, as we showed in \ref{QQQ}, we can usually assume this value is exceptionally small and omit it from the equation as we did in \ref{qlost}, \ref{qhot}, \ref{complete}. In this lab, we will be investigating the endothermic and exothermic qualities of salt solutions by dissolving various salts into water and monitoring the temperature. Then you will calibrate a coffee cup calorimeter by determining its heat capacity. Finally, you will calculate the specific heat of copper by dropping hot pennies into water and monitoring the heat exchange.

 

Experimental Procedure

Materials and Equipment

Sodium chloride, potassium chloride, calcium chloride, 10-15 pennies, DI water bottle. Equipment: 600 mL beaker, 250 mL beaker, 100 mL beaker, 50 mL graduated cylinder, 4- stryrofoam cups, 2 lids with hole, stirring rod, 3 small test tubes, thermometer, hot plate.

 

Safety

It is important to follow the safety guidelines below while performing this lab. Remember to dry the calorimeter and thermometer between each trial. Use appropriate procedures for hot plate use. Dispose of waste in the Waste container as indicated by your instructor. WEAR SAFETY GOGGLES, CLOSED-TOE SHOES.

 

Part A: Exothermic and Endothermic Dissolution of Salts

In this section of the procedure, you will observe temperature changes as various salts are dissolved in water.

1. The first salt is NaCl, and the corresponding dissolution reaction is \( \ce{NaCl(s)→Na+ (aq)+Cl− (aq)} \)
2. Record the mass of the empty test tube prior to beginning.
3. You will be recording the temperature using a thermometer. Insert the end of the thermometer into the bottom of a clean, dry test tube. (Use the small test tubes).
4. Fill the test tube approximately 2 cm with distilled water. Record the mass of the test tube and water to determine the mass of water.
5. Record the initial temperature of the water.
6. Record the mass of a second (clean and dry) test tube.
7. Fill the second test tube approximately 1 cm with solid NaCl.
8. Pour the solid NaCl into the water and stir gently with the thermometer.
9. Monitor the temperature. The temperature will increase or decrease away from the initial temperature. Eventually the final temperature will be reached before it begins to return to room temperature. Record the final temperature.
10. The solution should be disposed of in the waste container in the fume hood. Rinse and dry the test tube.
11. Repeat Steps 2-7 for CaCl2 and (if there is time) KCl.

 

Part B: Calculating the Heat Capacity of a Calorimeter

1. Obtain a hot plate and plug it in.
2. Stack the two Styrofoam cups together and place inside a 400-mL beaker
3. Place the assembly on a balance. Record the mass.
4. Place 50-mL of tap water in the cup assembly and record mass. Subtract to find the mass of water.
5. Place the cardboard lid on top the Styrofoam cups.
6. Insert a thermometer through hole in lid. Obtain the initial temperature of the “cold” water.
7. Attach a thermometer clamp to the thermometer so that the thermometer is not sitting on bottom of cup.
8. This is our calorimeter. 
9. Tare a 150mL beaker.
10. Add ~50mL of water to beaker on scale and record weight as the weight of “hot” water.
11. Place the ~150 mL beaker on the hot plate and heat to ~90 degrees.
12. Using the thermometer, record the temperature of the hot water.
13. Tip the lid of the calorimeter up and using beaker tongs immediately pour the hot water into the calorimeter.
14. Immediately replace the lid and begin recording temperature.
15. Gently swirl the calorimeter and measure temperature every 10 seconds until the temperature is constant for 3 readings. (The temperature will increase to the final temperature before beginning to decrease back to room temperature.) Record the final temperature in the data section.
16. Repeat steps 9-14 for a second trial.

 

Part C: Calculating the Specific Heat of Copper

1. Obtain the mass of ~ 10-20 copper shots or the weight of a small piece of copper wire. Record the mass in your data table.
2. Add copper shots/copper wire to a large (clean and dry) test tube. Place the test tube in a 400-mL beaker containing ~ 150 mL water. The water level should be above the level of the copper shots/copper wire to ensure they are adequately heated.
3. Heat water to ~ 95-100 o C. Record this as the initial temperature of the copper shots/copper wire. Allow the copper shots/copper wire to heat for ~2-3 minutes.
4. Add your calorimeter to the balance and tare it. Add ~40 mL of water. Record the mass of the water in the table.
5. Set up your calorimeter with thermometer as described earlier.
6. Record the initial temperature of the water.
7. Swiftly but carefully remove the test tube from the hot water bath with a pair of test tube tongs. Dump the copper shots/copper wire into the calorimeter and immediately cover with the lid. *Be careful that the hot water on the outside of the test tube does not drip onto your hand or into the calorimeter!
8. Monitor the temperature by taking temperature readings every 10 seconds. When the temperature is consistent for 3 straight measurements (or the temperature begins to cool) record the final temperature of the copper shots/copper wire and water.
9. Repeat steps 1-8 for a second trial. (Be sure to dry the copper shost between trials).
10. Use the data from your table to calculate the specific heat of copper (copper shots/copper wire).