Calorimetry Post Lab PDF

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Calorimetry Post lab...

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Someone has stolen my Joules! Question of the Day {QOD}: Can the experimentally derived Cp allow you to reliably identify an unknown metal? Safety: Safety Glasses

Nitrile Gloves Lab Coat Chemistry Concepts:  Law of Conservation of Energy  Heat/Energy – endothermic vs. exothermic  Specific heat capacity Techniques (refer to lab manual):  General Safety  Lab Equipment and Techniques o General Techniques o Measuring mass and volume o Calorimetry

 Quantitative Skills o Significant Figures o Data Analysis & Manipulation

Materials:     

Calorimeter 250mL Beaker, 400mL Beaker Vernier LabQuest 2 Temperature Probes 50mL graduated cylinder

   

Balance Hot plate Tongs Unknown solid object

Introduction Throughout your study of chemistry, you have observed it as a look into the physical and chemical changes of matter. However, chemistry also studies how energy is involved in all of these processes. Energy can take various forms when it is absorbed or released in a chemical process such as light, sound, work, or heat. All forms of energy can be transformed into heat (abbreviated q) and therefore energy changes can be described in terms of the amount of heat released to (exothermic processes) or absorbed from (endothermic processes) the surroundings. Changes in heat are convenient because they can easily be monitored by measuring the temperature changes of the system. Recall, however, that temperature and heat are not the same thing: temperature is a measure of how hot or cold an object is resulting from the movement of molecules, whereas heat is a form of energy. Imagine you are trying to boil a single cup of water to make coffee or tea vs. a large pot of water for pasta. At boiling, both are at the same temperature, but it took significantly more heat for the large pot to reach that temperature. Calorimetry is the study of heat transfer during a chemical reaction or from a physical change. The devices used to study these changes are known as calorimeters. Calorimeters come in a variety of shapes and sizes, but generally they are well-insulated devices sealed off from an external environment and contain a reaction chamber, stirrer, and thermometer. In all cases, the calorimeter provides a means to monitor heat flow between a system and its surroundings. Typically, when conducting calorimetry experiments the system is defined as the chemical reaction or physical process being studied. To more easily define the surroundings, often it is assumed that the calorimeter is perfectly insulated from the outside universe. Therefore, the surroundings in terms of the chemical or physical process is simply the calorimeter and whatever else was inside not directly involved in the

process, often water. Now because of the assumed isolation of the calorimeter and the law of conservation of energy, the amount of heat lost (or gained) by the system must be equal to the heat gained (or lost) by the surroundings. Clearly the heat of the system and surroundings must be of opposite sign since energy cannot be created or destroyed. This can be represented by the equation: qsystem = –qsurroundings Recall that the surroundings include the calorimeter and typically water so this can be rewritten as: –qsystem = qwater + qcalorimeter In addition to heat and temperature, a third term needs to be discussed in the context of thermodynamics: specific heat. The specific heat of a substance is the energy required to change one gram of that substance by 1oC. The units of specific heat, often abbreviated c, are therefore J/(goC). In order to rearrange for energy (or heat, in Joules), one could then write energy (J) = specific heat (J/(goC)) x mass (g) x T (oC) where T refers to the change in temperature of the substance and is calculated as Tfinal – Tinitial. The specific heat of a substance is a constant for a given state (for example c water = 4.184 J/goC) so the amount of energy absorbed or released by a substance will vary proportionally with both the amount of substance you have or the amount by which the temperature needs to change. It should be noted that T and q can have negative values if a substance cools down and emits heat, but c can never be negative. A negative specific heat implies that a substance would just be spontaneously heating up forever. It makes no sense physically. Again, think back to the boiling water example. A large pot of water has a much larger mass of water so it will require much more energy to heat to boiling. Similarly, it will require more energy to heat water from room temperature to boiling then say from room temperature to 40oC. As a general equation, these relationships can be summarized as qsubstance = msubstance x csubstance x Tsubstance Because the calorimeter is made of various materials, it is less useful to analyze the heat absorbed or lost by the calorimeter in terms of specific heat, but rather in terms of heat capacity, C, the amount of energy it takes to raise the temperature of an object by 1oC expressed in J/oC. Therefore, the amount of heat gained or lost by a calorimeter is often written as: qcalorimeter = Ccalorimeter x Tcalorimeter When trying to determine the heat involved in a chemical or physical process, it is important to first know the heat capacity of the calorimeter (sometimes called the calorimeter constant) being used in order to account for the exchange of heat with the device itself. A common way to determine the calorimeter constant is to graphically analyze the changes in temperature when hot water is added to cool water in a calorimeter. Thermodynamically, the amount of heat the hot water (which makes up the system) loses will be equal to the amount of heat the surroundings (comprised of the cool water and the calorimeter) gain. Since the specific heat of water is a known constant and masses and temperatures are easily measured in the lab, the calorimeter constant can be fairly easily calculated. One assumption that needs to be made in order to do so is that the calorimeter and cool water are starting and ending at the same temperature and therefore have the same T. Mathematically, all of this can be seen by manipulating the previously provided equations: qsystem = –qsurroundings qsystem = –(qwater + qcalorimeter) = –qwater + (–qcalorimeter)

qsubstance = msubstance x csubstance x Tsubstance qcalorimeter = Ccalorimeter x Tcalorimeter [(mH2O, hot)(cH2O)Thot] = -[(mH2O, cool)(cH2O)Tcool] + (-(CcalorimeterTcool)) Ccalorimeter = {[(mH2O, hot)(cH2O)Thot] + [(mH2O, cool)(cH2O)Tcool]} / (-Tcool) Ccalorimeter = {qhot + qcool} / -Tcool It is important to be very careful with signs in thermodynamics problems. Remember that in this experiment the hot water will be losing heat (-q) and the cool water will be gaining heat (+q) so the two added together when calculating the calorimeter constant should be a much smaller number than either individually. Just as in specific heat, the calorimeter constant should never come out negative. Another area to be careful of, particularly in determining the calorimeter constant, is the correct determination of the initial and final temperatures. To do so most accurately it is often best to generate a graph with four distinct lines – a regression showing the cool water before addition, one showing the warm water before addition, one showing the two after mixing, and a vertical line representing the time of mixing. An example of what this might look like can be seen below:

To find the most accurate temperatures, the regression equations are used where the time of mixing is plugged in for x (in this case 300 seconds). The corresponding y values for the hot and cold lines give the initial temperatures of the hot and cold water while the y value for the mixture line gives the final temperature of everything since it is assumed that the hot water, cold water, and calorimeter all reach thermal equilibrium. Therefore, in this example, Ti, cool = 24.0oC (there is no x component to this line), Ti, hot = (-0.0172 x 300) + 45.0 = 39.84oC, and Tf = (-0.0013 x 300) + 30.5 = 30.11oC. Now with these temperatures determined graphically, as long as the experimenter obtained accurate masses of both the cool and warm water, he or she can now calculate the calorimeter constant.

Once the calorimeter constant is known, the calorimeter can then be used for its intended purpose. In today’s

lab, that purpose will be to calculate the specific heat of an unknown solid object in order to try to correctly identify it. Much of the process for calculating a specific heat is similar to what has been described already. Here, instead of the system being a sample of hot water, it will be a sample of hot solid of known mass. This hot solid will be added to a calorimeter with a known quantity of cool water and the corresponding temperature changes will allow the experimenter to calculate the specific heat of the object using an adjusted version of the previous equations: qsystem = –(qwater + qcalorimeter) = –qwater + (–qcalorimeter) qsubstance = msubstance x csubstance x Tsubstance qcalorimeter = Ccalorimeter x Tcalorimeter [(msolid)(csolid)Tsolid] = -[(mH2Ol)(cH2O)TH2O] + (-(CcalorimeterTH2O)) csolid = {-[(mH2Ol)(cH2O)TH2O] + (-(CcalorimeterTH2O))} / [(msolid) Tsolid] csolid = -{qH2O + qcalorimeterl} / [(msolid)Tsolid] Ideally, temperatures for specific heat calculations would also be determined graphically, but since the specific heat of metals and metalloids is often significantly lower than that of water, and in the interest of time, we will simply approximate the initial temperature of the metal and the water as their respective temperatures immediately before mixing and the final temperature of the metal, water, and calorimeter, as the hottest temperature registered inside the calorimeter after mixing. This may not be the most accurate way of analyzing specific heat, but should serve as a reasonable approximation. Procedure Part I – Density of Unknown 1. Obtain an unknown solid sample from the supply cart and write down which unknown letter you take in your data table 2. Find the mass of your unknown solid and record it in the Table 1 3. Fill the smallest possible graduated cylinder that will hold your sample approximately halfway with tap water, recording the exact volume added 4. Carefully slide the solid into the cylinder at an angle, taking care not to shatter the glassware or spill the water 5. Record the combined volume of the solid and water in Table 1 6. Determine the volume of the solid and use it to calculate its density Part II – Calorimeter Constant Determination 7. Heat ~50 mL of tap water on a hot plate over medium-high heat to approximately 50oC, then remove the beaker from your hot plate and place it on a ceramic tile 8. Find and record the mass of your calorimeter in Table 3 9. Place ~50 mL of room temperature tap water into the calorimeter 10. Reweigh the calorimeter to find the exact mass of water inside (again, in Table 3) 11. Set up a LabQuest with two temperature probes, leaving one in the warm water on the ceramic tile and one in the cool water inside the calorimeter 12. When the temperature of the warm water reaches 45oC, begin recording the temperature of both the hot and cold water every 30 seconds over a 270 second period in Table 2 13. At the 300 second mark, remove the temperature probe from the hot water and quickly and carefully transfer it into the cold water in the calorimeter being careful not to spill

14. Constantly stir the combined contents and record the temperature of the mixture beginning at the 330 second mark until a final time of 600 seconds 15. While one person is stirring and recording temperatures, the other can begin heating a second sample of ~50mL of tap water for the second trial (just be careful not to heat this second sample too much or you will have to wait a while for it to cool) 16. After the 600 seconds is up, find the mass of the calorimeter to obtain the actual mass of hot water added and record it in Table 3 17. Repeat steps 12-16 one more time 18. Graph your data in Excel to determine the initial and final temperature values and then calculate the calorimeter constant (this step can be done at home) Part III – Specific Heat of an Unknown 19. 20. 21. 22.

Add about 250 mL of tap water to a 400-mL beaker and place it on a hotplate Add your unknown object and begin heating on high until the water boils While waiting for the water to boil, dry the calorimeter used in part II and find its mass Add enough tap water to your calorimeter to ensure the unknown object will be completely submerged once it is transferred 23. Reweigh the calorimeter to find the exact mass of water added 24. Leave the object in the boiling water for at least 5 minutes to ensure the unknown and water are at the same temperature 25. Find and record the temperature of the boiling water – you are assuming this is the initial temperature of your unknown 26. Find and record the temperature of the water in the calorimeter 27. Using tongs, carefully remove the solid object from the boiling water, gently shake it to remove excess water, and quickly transfer it to the calorimeter 28. Stir constantly and record the final temperature after 60 seconds 29. Repeat steps 18-27 two more times 30. Using the calorimeter constant from Part II and the data obtained in Part III, calculate the specific heat of your unknown (steps 30-31 can be done at home) 31. Try to identify which unknown you had based on the specific heat and density using Table 5 provided in the post-lab

Data Tables: Turn in with the post-lab report for this experiment. *Show a calculation for all marked boxes* Unknown Letter: ____G___ Table 1. Density for Unknown Mass Unknown (g)

4.434

Volume Water (mL)

15

Volume Water + Unknown (mL)

16

Volume Unknown (mL)*

1

Density Unknown (g/mL)*

4.434

Table 2. Calorimeter Constant Temperature Data Trial 1 Temperature Cold Water (oC)

Trial 2 Temperature Mixture (oC)

Temperature Hot Water (oC)

Temperature Cold Water (oC)

Time (s)

Temperature Hot Water (oC)

30

43.9

20.8

43.8

21.2

60

43.1

20.8

43.1

21.2

90

42.2

20.8

42.3

21.2

120

41.4

20.8

41.6

21.2

150

40.7

20.8

40.9

21.2

180

40.0

20.8

40.2

21.2

210

39.3

20.8

39.6

21.2

240

38.7

20.8

39.0

21.2

270

38.1

20.8

38.4

21.2

300

MIX

MIX

MIX

MIX

MIX

Temperature Mixture (oC)

MIX

330

27.3

28.1

360

27.3

28.1

390

27.3

28.1

420

27.2

28.1

450

27.2

28.1

480

27.2

28.0

510

27.1

28.1

540

27.1

28.0

570

27.1

28.0

600

27.0

27.9

Table 3. Calorimeter Constant Determination (Fill in the table below using the graphs generated from the time and temperature experimental data found in Table 2 – Using temperatures not found this way will result in significant loss of points!)

Trial 1 144.810

144.921

198.271

194.984

53.461

50.063

Tf mixture (oC)*

27.027

27.939

Ti cold (oC)*

20.8

21.2

T cold (oC)*

6.277

6.739

1404.04

1411.58

240.847

241.450

42.576

46.516

43.9

44.381

-16.873

-16.442

-3005.72

-3199.9

C calorimeter (J/oC)*

255.16

265.36

Average Ccal (J/oC)*

260.26



Mass Calorimeter (g)



Mass Calorimeter + Cool Water (g) Mass Cool Water (g)*

q cold (J)* 

Trial 2

Mass Calorimeter + Mixture (g) Mass Warm Water (g)* Ti hot (oC)* T hot (oC)* q hot (J)*

(  means data obtained in lab procedure)

Trial 1

Trial 2

Trial 3



Mass calorimeter (g)

144.705

145.548



Mass calorimeter + water (g)

278.242

261.517

Mass water (g)* 

Ti water (oC)



Tf (oC) T water (oC)* q water (J)*



Mass of unknown (g)



Ti unknown (oC)

236.201

133.537

115.969

91.54

21.1

21.4

21.7

21.4

21.6

22.1

0.3

0.2

0.4

167.62

97.04

153.20

78.078

52.052

104.1

4.434

4.457

4.446

99.3

99.5

99.8

-77.9

-77.9

-77.7

0.7101

0.429

0.745

260.26

Average Ccal (from Table 3) q calorimeter (J)*

144.661

T unknown (oC)* c unknown (J/(goC))* Average c unknown (J/(goC))* Table 4. Specific heat determination

(  means data obtained in lab procedure)

0.628

1. For Lab 8 (Calorimetry), attach all of the following: a. A properly formatted graph of your regression data for at least one trial of the calorimeter constant determination. (10pts)

b. All data tables completely filled out with sample calculations provided as designated (NOTE: CREDIT WILL NOT BE AWARDED WITHOUT THESE CALCULATIONS!) (30pts) Next page

’

c. Given the table below, determine the identity of your unknown object (be sure to provide the unknown’s ID letter). If you are nowhere near one of these options or close to several, list the possibilities. [If your technique was strong enough to correctly identify your object based on your data, you will earn 5pts extra credit on this post-lab]

The ID of the unknown metal was G. Based on my calculations, the specific heat of this metal was 0.628 J/g°C with a density of 4.434g/mL. Based on these numbers, we would be unable to identify which substance our metal could be. If trial 2 was an outlier, based on specific heat the unknown used in the experiment was silicon. Table 5: Specific Heat Capacity of various substances Substance

Specific Heat (J/goC)

Density (g/mL)

Substance

Specific Heat (J/goC)

Density (g/mL)

Aluminum

0.91

2.7

Lead

0.13

11.3

Bismuth

0.13

9.79

Magnesium

1.05

1.74

Brass

0.38

8.52

Nickel

0.44

8.90

Cadmium

0.23

8.69

Silicon

0.71

2.33

Chromium

0.46

7.15

Silver

0.24