Calorimeter Experiment

INTRODUCTION

Given appropriate chemicals and equipment, the specific heat capacity and molar mass of a metal, enthalpy of neutralization of an acid and base, and the enthalpy of solution of an unknown salt can be determined by following specific procedures. All of these procedures require the use of a calorimeter, which are of two types: a bomb calorimeter and a coffee cup calorimeter. Calorimeters are simply devices used to measure the amount of heat gained or lost in a system. Although this is not completely true, they are treated as isolated systems. A simple coffee-cup calorimeter can be constructed using two Styrofoam cups nested inside each other. The Styrofoam prevents heat loss to the surroundings, which makes it an ideal calorimeter.
The specific heat capacity of a substance refers to the amount of energy required to raise the temperature of one gram of substance by exactly 1°C. Given the specific heat capacity (c) of a substance, the mass (m), and the change in temperature (∆T), the amount of energy absorbed/released by a system can be calculated using the formula q = mc∆T [1]
If the specific heat capacity is known, the molar mass of the substance can be predicted using the following formula c x Mm = 25 J/mol °C.
It is important to note that the energy gained by the solution is equal to the energy lost by the surroundings, but opposite in sign. Therefore, qsolution = -qsurroundings [2]

The second part of the experiment was to calculate the enthalpy of neutralization, which follows the same concept as before. The energy gained by the solution is once again equal to the energy lost by the surroundings, but opposite in sign. To calculate the enthalpy of neutralization, the formula QN = -mc∆T can be used. However, once the enthalpy of neutralization is known, the amount of heat released per mole can be calculated using the formula
∆HN = QN / n [3]

The enthalpy of solution can be determined by performing an experiment in which a salt is dissolved into water. The temperature of the water is expected to decrease over time. By calculating the change in temperature and using the above formulas to calculate the energy gained by the salt and lost by the water, the enthalpy of solution can be calculated using the formula ∆Hs = -[ qwater / nsalt ] + [ - qsalt / nsalt ] [4]

By doing this experiment, the following results can be expected:
The specific heat capacity of the metal used (Zn) will be 0.39 J / g °C [1]
The molar mass of the metal used (Zn) will be 65.409 g / mol [2]
The heat of neutralization per mole of water for the acid HCl will be -57.9 kJ/mol [3]
The heat of neutralization per mole of water for the acid HNO3 will be -57.6 kJ/mol [3]

MATERIALS & PROCEDURE

As described in lab manual (What in the World ISN’T Chemistry?, Dr. Rashmi Venkateswaran, 2011, Exp. 3, p. 34-38). [4] OBSERVATIONS

Part 1: Specific Heat Capacity of a Metal

The distilled water used did not feel too hot or too cold to the touch. In fact, it was at room temperature, at approximately 23°C. The water was clean, so it was clear in color and had no apparent odour. It wasn’t very viscous; therefore, it flowed freely. The zinc metal was cut up into small rectangles. It was silver in color and quite brittle. It had no smell. It was in solid form; therefore, viscosity couldn’t be predicted.

Initially, when water was being heated over the hotplate, there was nothing to see. After a couple of minutes, bubbles started to appear as the water started to boil. Soon after, the water started to evaporate, and water vapour could be seen on the sides of the beaker. When the zinc was heated in a test tube, there seemed to be no change in its color or density. However, after it was added to the calorimeter and kept in there for several minutes, the color of the metal appeared to have changed from a shiny silver to a dull silver-gray color. It also appeared to be less dense, although it was still very hard. After the metal was added to the water in the calorimeter and the temperature was being monitored over 20 second intervals, it was noticed that each time the calorimeter was swirled, the temperature on the thermometer would increase and eventually return to its original state. This could have been due to the fact that when the cup was swirled, the needle of the thermometer touched the metal, which was much hotter than the water, causing the temperature to increase. Overall, the temperature of the solution at the end was higher than the temperature of the distilled water. At the end of the experiment, there seemed to be no apparent color change in the overall solution.

Part 2: Enthalpy of Neutralization

HCl and NaOH

In this experiment, HCl was the acid and NaOH was the base. Both substances were clear in color and seemed to be as viscous as the water in Part 1. The concentration of the HCl was 1.0M and the concentration of the NaOH was 1.1M. Initially, the base was at room temperature, at 22.6°C. When the acid was added to it, the overall temperature of the solution significantly rose to 29.5°C. It eventually came down to 29.0 °C.

Part 3: Enthalpy of Solution The distilled water was originally at room temperature, at 21.9 °C. It was colorless and odourless, and not very viscous. The salt was in the form of a very soft, white powder. It was also odourless. Initially, the distilled water was in the calorimeter, and its temperature was being monitored over 30-second intervals. It was noticed that when the salt was added, the temperature of the solution started to decrease. It decreased from 22.4 °C to 16.0 °C, which is a difference of 6.4 °C.
RESULTS

Part 1: Enthalpy of a Metal

Data
Trial 1
Identity of Metal
Zn
Change in Water Temperature
2.50 °C
Energy Gained by Water
198 J
Change in Metal Temperature
-72.9 °C
Experimental Specific Heat Capacity of Metal
0.197 J/g °C
Accepted Specific Heat Capacity of Metal
0.39 J/g °C
Experimental Molar Mass of Metal
127 g/mol
Accepted Molar Mass of Metal
65.409 g/mol
Percent Error (Specific Heat Capacity)
-49.5%
Percent Error (Molar Mass)
94.2%

Part 1: Sample Calculations (TRIAL 1)

1. From your graph, calculate the change in temperature of the water, ∆Twater.

∆Twater = Tfinal - Tinitial = 27.1 °C – 24.6 °C = 2.50 °C

2. Using the specific heat capacity of water and the mass of water, calculate the energy gained by the water.

c = 4.184 J/g °C
∆Twater = 2.50 °C

mwater = m(calorimeter & water) – m(empty calorimeter) = 27.9117 g – 8.9818 g = 18.9299 g qwater = (mwater) (c) (∆Twater) = (18.9299 g) (4.184 J/g °C) (2.50 °C) = 198 J
3. Calculate the change in temperature of the metal, ∆Tmetal.

∆Tmetal = Tfinal - Tinitial = 27.1 °C – 100 °C = -72.9 °C

4. Determine the specific heat capacity of the metal, using the energy gained by the water and the mass of the metal.

qmetal = 198 J mmetal = 13.76 g
∆Tmetal = -72.9 °C

cmetal = -qmetal (mmetal) (∆Tmetal) = 198 J (13.76 g) (-72.9 °C) = 0.197 J/g °C

5. Approximate the molar mass of the metal using its calculated specific heat capacity.

cmetal = 0.197 J/g °C

Mmmetal = 25 J/mol °C cmetal = 25 J/mol °C 0.197 J/g °C = 127 g/mol

6. Since you know the identity of the metal, calculate a percent error with respect to the known specific heat capacity AND known molar mass of the metal.

cexperimental = 0.197 J/g °C cactual = 0.390 J/g °C

% error = cexperimental - cactual x 100% cactual = (0.197 J/g °C) – (0.390 J/g °C) x 100% 0.390 J/g °C = -49.5%

Mmexperimental = 127 g/mol Mmactual = 65.39 g/mol

% error = Mmexperimental - Mmactual x 100% Mmactual = (127 g/mol) – (65.39 g/mol) x 100% 65.39 g/mol = 94.2%

Part 2.1: Enthalpy of Neutralization

Data
Trial 1
Identity of Acid
HCl
Change in Temperature
6.40 °C
Volume of Final Solution
100. mL
Mass of Final Solution
102.5 g
Energy Released
-2.74 kJ
Number of Moles of Limiting Reagent (OH-)
0.050 mol
Number of Moles of Water Formed
0.050 mol
Experimental Heat of Neutralization / Mole of Water
-55 kJ/mol
Accepted Heat of Neutralization / Mole of Water
-57.9 kJ/mol
Percent Error
-4.51%

Part 2.1: Sample Calculations (TRIAL 1)

7. Determine the change in temperature of the solution, ∆Tsoln.

∆Tsoln = Tfinal - Tinitial
= 29.0 °C – 22.6 °C = 6.40 °C

8. Calculate the volume of the final solution.

Vfinal = VHCl + VNaOH = 50.00 mL + 50.00 mL = 100. mL

9. Calculate the mass of the final solution, assuming the final solution to have a density of approximately 1.0 g/mL. How does this compare to your measured mass? Which is more accurate?

d = 1.0 g/mL
V = 100.0 mL m = dV (mass according to density) = (1.0 g/mL) (100. mL) = 100. g

m = mcalorimeter & solution – mcalorimeter (measured mass) = 111.44 g – 8.9618 g = 102.5 g

Measured mass is more accurate because in the calculated mass, the density is being assumed. If the actual density was known, the calculated mass would have been more accurate. From this point on, the measured mass will be used in the calculations.

10. Calculate the energy released, assuming that the specific heat of the final solution is the same as that of water, 4.184 J/g °C.

c = 4.184 J/g °C m = 102.5 g
∆T = 6.4 °C

q = -mc∆T = -(102.5 g) (4.184 J/g °C) (6.4 °C) = -2.74 x 10 3 J = -2.74 kJ

11. Calculate the number of moles of the limiting reagent (OH-(aq)).

cNaOH = 1.1 M
VNaOH = 50.00 mL = 0.050 L

nNaOH = (cNaOH) (VNaOH) = (1.0 M) (0.050 L) = 0.050 mol NaOH

nOH = nNaOH (because Na to OH ratio = 1:1) = 0.050 mol OH

12. Calculate the number of moles of water formed in the neutralization reaction.

HNO3 + NaOH  NaNO3 + H2O

nH2O = 0.050 mol NaOH x 1 mol H2O 1 mol NaOH = 0.050 mol H2O

13. Determine the heat of neutralization per mole of water.

nH2O = 0.050 mol QN = -2.74 kJ

∆HN = QN nH2O = -2.74 kJ 0.050 mol = -55 kJ/mol

14. Compare the heats of neutralization per mole of water for the two strong acids.
This section is answered in the discussion.

15. Calculate the percent errors for the experimental values of the heat of neutralization with respect to literature values.

∆HN-experimental = -55 kJ / mol ∆HN-actual = -57.9 kJ/mol

% Error = ∆HN-experimental - ∆HN-actual x 100% ∆HN-actual = (-55 kJ / mol) – (-57.9 kJ/mol) x 100% -55 kJ/mol = -4.5%

Part 2.2: Enthalpy of Neutralization

Data
Trial 1
Identity of Acid
HNO3
Change in Temperature
6.50 °C
Volume of Final Solution
100. mL
Mass of Final Solution
101.4 g
Energy Released
-2.76 Kj
Number of Moles of Limiting Reagent (OH-)
0.050 mol
Number of Moles of Water Formed
0.050 mol
Experimental Heat of Neutralization / Mole of Water
-55 kJ/mol
Accepted Heat of Neutralization / Mole of Water
-57.6 kJ/mol
Percent Error
4.51%

Part 2.2: Sample Calculations (TRIAL 1)

7. Determine the change in temperature of the solution, ∆Tsoln.

∆Tsoln = Tfinal - Tinitial
= 31.6 °C – 25.1 °C = 6.50 °C

8. Calculate the volume of the final solution.

Vfinal = VHNO3 + VNaOH = 50.00 mL + 50.00 mL = 100. mL

9. Calculate the mass of the final solution, assuming the final solution to have a density of approximately 1.0 g/mL. How does this compare to your measured mass? Which is more accurate?

d = 1.0 g/mL
V = 100. mL

m = dV (mass according to density) = (1.0 g/mL) (100. mL) = 100. g

m = mcalorimeter & solution – mcalorimeter (measured mass) = 110.3709 g – 8.986 g = 101.4 g

Measured mass is more accurate because in the calculated mass, the density is being assumed. If the actual density was known, the calculated mass would have been more accurate. From this point on, the measured mass will be used in the calculations.

10. Calculate the energy released, assuming that the specific heat of the final solution is the same as that of water, 4.184 J/g °C.

c = 4.184 J/g °C m = 101.4 g
∆T = 6.50 °C

q = -mc∆T = -(101.4 g) (4.184 J/g °C) (6.5 °C) = -2.76 x 10 3 J = -2.76 kJ

11. Calculate the number of moles of the limiting reagent (OH-(aq)).

cNaOH = 1.0 M
VNaOH = 50.00 mL = 0.050 L

nNaOH = (cNaOH) (VNaOH) = (1.0 M) (0.050 L) = 0.050 mol NaOH

nOH = nNaOH (because Na to OH ratio = 1:1) = 0.050 mol OH

12. Calculate the number of moles of water formed in the neutralization reaction.

HNO3 + NaOH  NaNO3 + H2O

nH2O = 0.050 mol NaOH x 1 mol H2O 1 mol NaOH = 0.050 mol H2O

13. Determine the heat of neutralization per mole of water.

nH2O = 0.050 mol QN = -2.76 kJ

∆HN = QN nH2O = -2.76 kJ 0.050 mol = -55 kJ/mol

14. Compare the heats of neutralization per mole of water for the two strong acids.
This section is answered in the discussion.

15. Calculate the percent errors for the experimental values of the heat of neutralization with respect to literature values.

∆HN-experimental = -55 kJ / mol ∆HN-actual = -57.6 kJ/mol

% Error = ∆HN-experimental - ∆HN-actual x 100% ∆HN-actual = (-55 kJ / mol) – (-57.6 kJ/mol) x 100% -57.6 kJ / mol = 4.5%
Part 3: Enthalpy of Dissolution of a Salt

Data
Trial 1
Identity of Salt
Unknown Salt B / Potassium chloride (KCl)
Change in Temperature of Solution
-5.20 °C
Energy Absorbed/Released by Water
-204 J
Energy Absorbed/Released by Salt
-30.0 J
Experimental Enthalpy of Dissolution
15.9 kJ/mol
Accepted Enthalpy of Dissolution
17.2 kJ/mol
Percent Error
-7.56 %

Part 3: Sample Calculations (TRIAL 1)

16. Determine the change in temperature of the solution, ∆Tsoln.

∆Tsoln = Tfinal - Tinitial = 16.0°C – 21.2 °C = -5.20 °C

17. Calculate the energy released/absorbed by the solution, using its specific heat and mass.

mwater = 9.3826 g cwater = 4.184 J/g °C ∆Tsoln = -5.20 °C

qwater = (mwater)( cwater) (∆Tsoln) = (9.3826 g) (4.184 J/g °C) (-5.20 °C) = -204 J

msalt = 1.49 g csalt = 3.877 J/g °C ∆Tsoln = -5.20 °C

qsalt = (msalt)( csalt) (∆Tsoln) = (1.49 g) (3.877 J/g °C) (-5.20 °C) = -30.0 J

18. Calculate the enthalpy of dissolution, ∆Hs, per mole of salt.

msalt = 1.49 g Msalt = 101.11 g/mol nsalt = 0.0147 mol

qwater = -204 J qsalt = -30.0 J

∆Hs = -[ qwater / nsalt ] + [ - qsalt / nsalt ] = - [-204 J / 0.0147 mol] – [-30.0 J / 0.0147 mol] = 1.59 x 10 4 J/mol = 15.9 kJ/mol

19. Attempt to find a literature value for the enthalpy of dissolution of the salt (but make SURE you are looking for the correct quantity!) If you find a literature value, calculate the percent error.

Literature Value for ∆Hs = 17.2 kJ/mol (Potassium chloride (KCl)) [6] Experimental Value for ∆Hs = 15.9 kJ/mol

% Error = (Experimental Value) – (Literature Value) * 100% (Literature Value) = (15.9 kJ/mol) – (17.2 kJ/mol) * 100% (17.2 kJ/mol) = -7.56 %

DISCUSSION

Part 1: Specific Heat Capacity of a Metal In this part of the experiment, 13.76 g of zinc metal at 100.0 °C was added to 20.00 mL of distilled water at room temperature.
This part of the experiment yielded an exothermic reaction. The purpose was to calculate the specific heat capacity and molar mass of the zinc metal. The specific heat capacity was found to be 0.197 J/g °C with a percent error of -49.5%. The molar mass was found to be 127 g/mol with a percent error of 94.2%. Possible sources of error inherent to this segment of the experiment were the fact that when the temperature of the solution was being measured, the calorimeter was placed near the hotplate by mistake. The heat from the hotplate could have transferred to the calorimeter, which would result in an increased temperature. This explains why the percent error for the specific heat capacity was negative. Another source of error is that the hole in the lid of the calorimeter gradually expanded, which resulted in a heat loss from the calorimeter. A second source of error is that the graduated cylinder used at the beginning of the experiment was not clean. The dust particles in the graduated cylinder could have reacted with the water or with the metal, resulting in a change of mass of the overall solution. This would result in the specific heat capacity of zinc being altered, due to the fact that the dust particles most likely have a specific heat capacity different than that of
zinc.

Part 2: Enthalpy of Neutralization In this part of the experiment, two different acids were reacted with the base NaOH to form a salt.
These reactions were exothermic. The purpose of this part of the experiment was to calculate the heat of neutralization per mole of water. Two trials were conducted using HCl to form NaCl salt. The heat of neutralization per mole of water was found to be -55 kJ/mol with a percent error of -4.51%. Two trials were conducted using HNO3 to form NaNO3 salt by the two students mentioned in the references section. The heats of neutralization per mole of water were found to be -57.6 kJ/mol, with a percent error of -4.51%. Although the percent errors weren’t very large, they were still off. These could have occurred due to the fact that not all of the acid solution was poured into the calorimeter. Some of the acid could have reacted with the air, and this would decrease the volume of acid which actually mixed with the base, resulting in an increased heat of neutralization. A second source of error was that in the experiment where HCl and NaOH were being mixed, one drop of NaOH was lost by accident which the NaOH solution was being poured into the calorimeter. This would result in too much acid solution, which means that not all of the acid completely reacted with the base. This, also, would increase the heat of
neutralization.

Part 3: Enthalpy of Solution In this part of the experiment, an unknown salt was added to distilled water and the temperature of the new solution was monitored over 20-second intervals. The purpose of this part of the experiment was to calculate the enthalpy of dissolution of the salt. The trials were conducted using an Unknown Salt B, which is assumed to be KCl. The enthalpy of dissolution was calculated to be 15.9 kJ/mol with a percent error of -7.56%. The negative percent error can be explained by addressing the fact that not all of the salt was actually added to the distilled water. What happened was that in order to measure the exact mass of the salt, it was put on a weighing paper in order to put on an analytical balance. When the salt was being poured from the weighing paper into the calorimeter, some of the salt crystals would not detach from the weighing paper. This resulted in less amount of salt dissolving with the water, which results in an enthalpy of dissolution lower than the actual enthalpy of dissolution, thus the negative percent error. One noticeable factor in this part of the experiment, which was different from the other parts was the fact that after the salt was added, the temperature of the solution decreased instead of increasing. This can be explained by simply considering the hydrogen bonding between water molecules. In order to dissolve, the salt needed energy, which is gained through disrupting the hydrogen bonds between water molecules. By disrupting the hydrogen bonds between water molecules, the salt gained energy which helped it dissolve. The fact that hydrogen bonds were disrupted resulted in an endothermic reaction, which explains why the temperature of the solution decreased.

CONCLUSION

The average specific heat capacity and molar mass of zinc was calculated to be 0.197 J/g °C and 127 g/mol, with average percent errors of -49.5% and 94.2% respectively. The average heat of neutralization per mole of water for the reaction between NaOH and HCl was calculated to be -55 kJ/mol, with an average percent error of -4.51%. The average heat of neutralization per mole of water for the reaction between NaOH and HNO3 was calculated to be -55 kJ/mol, with an average percent error of 4.51%. Last but not least, the average enthalpy of dissolution of Unknown Salt B, now assumed to be potassium chloride, was calculated to be 15.9 kJ/mol, with an average percent error of -7.56%. These values were in proximity to the accepted values.

REFERENCES

[1] http://www2.ucdsb.on.ca/tiss/stretton/database/Specific_Heat_Capacity_Table.html

[2] http://www.lenntech.com/periodic/mass/atomic-mass.htm

[3] David R. Lide, ed., CRC Handbook of Chemistry and Physics, 89th Edition (Internet Version 2009), CRC Press/Taylor and Francis, Boca Raton, FL

[4] Rashmi Venkateswaran, 2011, Experiment 3: Enthalpy of Various Reactions, What in The
World ISN’T Chemistry? General Chemistry CHM 1301/1311 2011, p.34-38, Ottawa Ontario, Canada, University of Ottawa/Faculty of Sciences

[5] Raw data table provided by CHM1311 students Vanessa Sinden-Lafleche and William Maich who were supervised by TA Daniela Marquez of the University of Ottawa/Faculty of Sciences

[6] http://www.mindspring.com/~drwolfe/WWWolfe_dat_enthalpies.htm