Construction of Calorimeters and Calculation of Heat of Reaction

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Construction of Calorimeters and Calculation of Change in Enthalpy of Acid-

Base Reaction

Santiago Charry, Christopher Stanley, Abigail Harrison, Micah Perwein

Abstract: The purpose of the lab was to construct two different calorimeters, calculate their heat

capacities, use them to find the change in enthalpy of the acid-base reaction between NaOH and

HCl, and find if a change in volume of the acid and/or base had any effect on the change in

enthalpy. The calorimeters were assembled using a variety of materials including Styrofoam

cups, duct tape, cardboard, and thermometers. The heat capacities were discovered by recording

changes in water temperature when hot water was added to cold water in the calorimeters. The

change in enthalpy of the acid-base reaction between NaOH and HCl was found by recording

temperature changes when the substances were reacted within the calorimeters and using

enthalpy equations to calculate the change in enthalpy of the reaction. A variation in the volumes

of each substance was tested over a series of trials to determine if it had any effect on the change

in enthalpy of the reaction. The results of the experiments indicated that our first calorimeter had

a heat capacity of 57.53 J/C, our second calorimeter had a heat capacity of 14.46 J/C, and the

enthalpy of the reaction between NaOH and HCl was -75.63 KJ/mol. The variation in volume of

each substance did not appear to have an effect on the change in enthalpy per mole of product.

Introduction

Two different calorimeters were assembled using a variety of household materials. After

their construction, their heat capacities were calculated. The calorimeters were then used to find

the enthalpy of the acid-reaction between NaOH and HCl and observe correlation between

volume of substances and change in enthalpy. The ability to construct and use calorimeters to


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find enthalpies of chemical reactions is often applicable to scientists doing lab work. According

to J. Phillipson from the Department of Zoology at the University of Durham, ecologists

interested in population and community energetics have been stressing the need for a more

complete knowledge of the energy content of biological materials.1. The ability to construct

different types of calorimeters and find enthalpies of reactions might enable ecologists to find the

energy content of such biological materials. The ability to calculate the enthalpy of reactions also

proved useful for scientists analyzing the behavior of antigens and antibodies. By observing that

the enthalpy of reactions was proportional to a decrease in water activity in the reactions of

lysozyme with a specific antibody, the scientists were able to provide further support that water

molecules bound to the antigen and antibodies are conserved upon formation for stability.2In

this report, readers can expect to learn how to construct effective calorimeters out of common

materials, how to calculate their heat capacities, and how to use them to find the change in

enthalpy of acid-base reactions. In doing this, they will gain an understanding of how to create an

effective calorimeter to insulate heat, and learn how to find the energy content of different

reactions.

Experimental

Construction of Calor imeters

Calorimeter #1: The construction of the first calorimeter required the use of two Styrofoam cups,

each with a height of 8.4 cm, a diameter of 4.4 cm at the bottom of the cup, and a diameter of 7.5

cm at the top of the cup. In addition, it required 16.1 mL of water, duct tape of the same length as

the circumference of the top of the cup, and two circles of cardboard with the same

circumference as the cups opening. To assemble the calorimeter, one of the Styrofoam cups was

filled with 16.1 mL of water. The second cup was then inserted into the first cup, deep enough


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that the water level of the first cup reached its rim without overflowing. Duct tape was then

wrapped around the length of the first cups rim against the side of the second cup once, sealing

the water in. Two circles with the same circumference as the cups opening were cut out from

cardboard. These circles were stacked on top of each other, creating a double-layer of cardboard,

and duct taped together along their edges. A hole large enough to provide a snug fit for our

thermometer to enter the calorimeter was poked into the double layer of cardboard. The

cardboard circle was then inserted into the opening of the Styrofoam cup, completing our

calorimeter.

Calorimeter #2:The construction of the second calorimeter involved three Styrofoam cups of the

same dimensions as those of the first calorimeter, 32.2 mL of water, two pieces of duct tape of

the same length as the circumference of the cups rim, and a Styrofoam circle with the same

circumference as the cups opening. To assemble this calorimeter, one of the cups was filled with

16.1mL of water. A second cup was inserted into the first, deep enough that the water level of

the first cup reached its rim without overflowing. The rim of this cup was sealed against the side

of the second cup using a strip of duct tape with the same length as the rims circumference. The

second cup was then also filled with 16.1 mL of water. The third Styrofoam cup was inserted

into the second cup, deep enough that the water level of the second cup reached its rim without

over flowing. The rim of the second cup was sealed against the side of the third cup using a strip

of duct tape with the same length as the rims circumference, as previously done with the first

and second cups. A circle with the same circumference as the cups opening was then cut out

from a piece of Styrofoam. A hole large enough to provide a snug fit for our thermometer to

enter the calorimeter was poked into this circle of Styrofoam. The circle was inserted into the

opening of the third Styrofoam cup, completing our second calorimeter.


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Calculation of Heat Capacities of Calorimeters

The heat capacity of each calorimeter was found by recording changes in water temperature

when hot water was added to cold water within the calorimeters. 80 mL of water was measured

in a 100-mL graduated cylinder. The graduated cylinder was then inserted into a beaker filled

with ice in order to cool down the water. In the meantime, 80 mL of water was measured in a

different 100-mL graduated cylinder then poured into a 250-mL beaker. The beaker was placed

on a hot plate to heat this water. Caution was taken not to allow the water to heat past boiling

point, so as to prevent it from evaporating. Let it be noted that 85mL of hot and cold water was

used when to find the heat capacity for calorimeter #2. After time was given to allow for a

significant change in temperature, the temperatures of the hot and cold water were recorded. The

cold water was poured into the calorimeter, and the hot water was immediately added. The

highest stable temperature of the combined water was recorded. Calorimetric equations were

used with the newly obtained data to find the heat capacity of the calorimeter. This method was

repeated with the second calorimeter to identify its heat capacity.

Calculation of the Change in Enthalpy for Acid-Base Reaction Between NaOH and HCl

The calculation of the change in enthalpy for the acid-base reaction between NaOH and HCl was

done by recording temperature changes when the substances were reacted within the calorimeters

and using enthalpy equations to calculate the change in enthalpy of the reaction. A variation in

the volumes of the substances was added to observe correlation between volume of substances

and enthalpy of the reaction. This was done in a series of six trials, three of which taking place in

the first calorimeter, and the other three taking place in the second calorimeter. First, specific

amount of the reactants were obtained: 80 mL of 1M NaOH and 80 mL of 1M HCl for trials 1 &

2, 60 mL of 1M NaOH and 100 mL of HCl for trials 3 & 4, and 100 mL of 1M NaOH and 60 mL


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of 1M HCl for trials 5 & 6. The initial temperature of each substance was then taken. Both

reactant were added to the calorimeter (calorimeter 2 used for trials 1, 3, and 5; calorimeter 1

used for trials 2, 4, and 6). The highest stable temperature during the reaction within the

calorimeter was recorded, and used in enthalpy equations to find the change in enthalpy for the

acid-base reaction. Once all six trials were completed, the average change in enthalpy per mole

was calculated. Waste management throughout the lab involved throwing away unneeded scraps

during construction of calorimeters, neutralizing the acids and bases and pouring the neutralized

product in the drain, and rinsing glassware and equipment with deionized water and allowing it

to air dry.

Results and Discussion

Calculation of Heat Capacities of Calorimeters

The data table below (table 1) displays the recorded temperatures of the hot and cold water used

in the calculation the heat capacities of calorimeters 1 &2.

Table 1: Initial and Final Temperatures of Hot and Cold Water

Data for Calorimeter #1 Data for Calorimeter #2

Relative

Temperature of

Water

Initial

Temperature

Final

Temperature

Initial

Temperature

Final

Temperature

Cold Water 16.8C 48.8C 12.8C 49.7C

Hot Water 86.3C 48.8C 88.1C 49.7C

Since the cold water was added to the calorimeter first, and then had its temperature recorded,

the calorimeter was assumed to have the same initial and final temperature as the cold water.

These values were then used to calculate the heat capacity of the calorimeters. The heat capacity

(Ccal) of the calorimeters demonstrates the amount of heat in kilojoules that escapes the


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calorimeter for every change in degree of Celsius or Kelvin. The greater the heat capacity, the

more heat escapes per change in degree. The calculations for this are shown below.

Calculations for Calorimeter # 1

80ml = 80g

Qcold = 80g 4.184 J/(gK) 32K = 10711.04 J

Qhot = 80g 4.184 J/(gK) -37.5K = -12552.0 J

10711.04J + -12552.0J = -Ccal 32K

57.53 J/K = Ccal

Calculations for Calorimeter # 2

85ml = 85g

Qcold = 85g 4.184 J/(gK) 36.9K = 13123.12 J

Qhot = 85g 4.184 J/(gK) -38.4K = -13656.58 J

13123.12J + -13656.58J = -Ccal 36.9K

14.46 J/K = Ccal

A possible source of error in attaining these values is any escape in heat that may have occurred

when capping the calorimeters, however, such an escape in heat would likely have occurred for

the reaction between NaOH and HCl as well, since capping the calorimeters worked in the same

way for both experiments.

Calculation of the Change in Enthalpy for Acid-Base Reaction Between NaOH and HCl

The following table (table 2) displays the volume of the reactants used for each trial, the initial

temperatures of reactants NaOH and HCl, and the final temperatures of their produced

compound. It was assumed that the NaOH and HCl had the same initial temperature, since both

were being kept in the same environment.


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Table 2: Volume and Initial and Final Temperatures of NaOH and HCl

Trial # Volume of

1M NaOH

Volume of

1M HCl

Calorimeter

used

Initial

Temperature

Final

Temperature

1 80mL 80mL 2 23.6C 31.2C

2 80mL 80mL 1 23.8C 31.3C

3 60mL 100mL 2 25.8C 33.9C

4 60mL 100mL 1 23.4C 30.7C

5 100mL 60mL 2 23.3C 29.3C

6 100mL 60mL 1 23.4C 29.3C

Using this data, enthalpy equations could be used to find the change in enthalpy of the acid-base

reactions in each trial. These calculations are shown below. It is assumed that mL=g since the

reaction results in water, which has a density of 1mL/g, so the conversion from mL to gram is

not shown. Since the substances have a molality of 1, the moles of each substance are equal to

their liters, so this conversion factor is not shown in the calculations. Furthermore, since the

reaction has a 1:1 ratio between all substances in a balanced chemical equation

(NaOH+HClNaCl+H2O), enthalpy can be noted in units joule per mole of any substance

involved in the reaction. Since this 1:1 ratio causes whichever reactant with the lowest volume to

be the limiting reactant, then moles of substance reacted are always based on whichever reactant

has the lowest volume.

Trial #1

Qsolution = 4.184 J/(gK) 160g 7.6K = 5087.7 J

Qcal = 14.46 J/K 7.6K = 109.896 J

5087.7J + 109.896J = -Qrxn


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HCl fHo(kJ/mol) = -167.2

NaCl fHo(kJ/mol) = -407.27

H2O fHo(kJ/mol) = -285.82

NaOH + HCl NaCl + H2O

(-407.27 + -285.82)(-469.15 + -167.2) = -56.77 KJ/mol Hrxn

Shown below are calculations for percent error of our experimental calculation of the change in

enthalpy.

(-75.63 experimental value - -56.77 theoretical value) / -56.77 theoretical value = .33 = 33%

error

According to the calculations, our experiment resulted in a 33% error in the calculated change in

enthalpy. A possible source of error in these calculations may be in the heat capacity of the

calorimeters. The calorimeters were allowed to sit for seven days after the measurement of heat

capacity. Upon return, the insulation of the calorimeters, which were supposed to hold 16.1mL of

insulating water between the Styrofoam cups tightly enough that they were silent if shaken, now

made a swishing noise if shaken lightly, suggesting that there was now air as well as water

between the Styrofoam cups of the calorimeter. This may have changed the experimental heat

capacity to a value different than what was used for calculation, and resulted in skewed data.

Another possible error that may have contributed to the 33% error is the temperature of the

substances themselves. There is the possibility that the graduated cylinders they were collected in

had been washed in water of differing temperature, which may have contributed to a temperature

change caused not by the reaction, but by the initial substance of each substance itself. Since we

only recorded the initial temperature of one of the reactants, if the other reactant was a different


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temperature it may have contributed to the change in temperature.

Conclusion

According to the testing done in the lab, the first calorimeter constructed had a heat capacity of

57.53 J/C and the second had a heat capacity of 14.46 J/C. The change in enthalpy of the acid-

base reaction between NaOH and HCl was -75.63 KJ/mol with a 33% error, and there was no

clear trend between volume of substance reacted and enthalpy of reaction, although such a trend

may have emerged with further testing of more drastic volume changes. Also, although we

varied the amount of each reactant for the six trials, four of the trials used the same amount of

limiting reactant, which would not differ the amount of change in enthalpy. This would be

another reason that the trend between volume of substance and enthalpy of reaction was unclear.

The construction and application of calorimeters as exemplified in this lab is a significant skill

for chemists to develop. Such a skill can prove useful for chemists even in the field of food

science, in which determining the amount of calories in food is necessary for labeling so that

consumers can properly monitor their diet.3Using calorimeters to find enthalpy changes can be

applied to many fields of science and is therefore significant for scientists to understand.

References

(1) Phillipson, J. A Miniature Bomb Calorimeter for Small Biological Samples

Oikos2007, 15, 130139.

(2) Goldbaum, F. A.; Schwarz, F. P.; Eisenstein, E.; Cauerhff, A.; Mariuzza, R. A.; Poljak,

R. J. The Effect of Water Activity on the Association Constant and the Enthalpy of Reaction

Between Lysozyme and the Specific Antibodies D1.3 and D44.1.J. Mol. Recognit.9, 612.

(3) Miller, D. S.; Payne, P. R. A Ballistic Bomb Calorimeter.Br. J. Nutr.1959, 13, 501508.

Construction of Calorimeters and Calculation of Heat of Reaction

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