Cortez Rankin
Team II
College of Engineering
University of Tennessee at Chattanooga
615 McCallie Avenue
Chattanooga, TN 37403
To:
Dr. Frank Jones.
Professor of Engineering
University of Tennessee at Chattanooga
615 McCallie Avenue
Chattanooga, TN 37403
Dr. Henry:
The following report provides a an account of Team II’s experimental work and all other
relevant information related to the study of the heats of solution experiment of
ammonium nitrate in water done in EMCS 120 at the University of Tennessee at
Chattanooga located in.
Cortez Rankin
Team II
Senior Undergraduate Student
Chemical Engineering
University of Tennessee at Chattanooga
HEATS OF SOLUTION EXPERIMENT
University of Tennessee at Chattanooga
College of Engineering
Engineering 435 Chemical Process Laboratory
By: Cortez Rankin
To: Dr. Frank Jones
December 7, 2003
TABLE OF CONTENTS
Page
Introduction…………………………………………………………………………………1
Theory………………………………………………………………………………………2
Equipment…………………………………………………………………………………..4
Procedure……………………………………………………………………………………5
Results………………………………………………………………………………………6
Discussion of Results………………………………………………………………………..11
Conclusions………………………………………………………………………………….12
References…………………………………………………………………………………...13
INTRODUCTION
The purpose of this experiment was for Team II to determine the heats of solution for
ammonium nitrate in water. Then, Team II was to compare the values calculated from
this experiment with the listed values located in a chemistry handbook. The experiment
began by taking a 5mL vial of ammonium nitrate (8.65g) and mixing the contents into
various amounts of water ranging from 20 to 240mL. From there, the change in
temperature was determined, and the heats of solution were calculated for the various
solution compositions. The heats of solutions were measured at 20mL intervals.
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THEORY
When a solid dissolves in water, the process always has energy change associated with
it. Examples exist for both endothermic and exothermic heats of solution. However, the
dissolving process itself is really a two-step process. The first step, breaking down the
solid crystal, is endothermic while the second step, hydrating the individual particles
released into the solvent, is exothermic. The overall heat of solution depends on the
relative amounts of energy involved in the two individual steps. With the addition of
ammonium nitrate into a water solution, this two-step process is analyzed thus
determining whether an exothermic or endothermic reaction is taking place.
Start
W0
Q 0
H  0
End
(1) A  N  H2 O  A  N  H2 O
( s)
(l )
( solution )
A-N in equation 1 represents ammonium nitrate in its solid state. H2O represents water.
This equation states that no heat or work was added to the solution to drive this reaction.
@18 C
@18 C
H
 (  H s )n A N
s ,18 C
 A  N  H2 O
(2) A  N  H2 O 
( s)

(l )
( solution )
Equation two shows the same reaction, but shows the heat of solution change that occurs
at 18oC.
H
s ,18 C
 (  H s )n A  N
shows the amount of heat of solution that it requires to
drive this reaction. nA-N represents the number of moles of ammonium nitrate.
(3)  H  nC p  T
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The heats of solution had to be calculated using the change in energy from the
experiment that was needed to obtain the values at 18oC. T is the temperature
difference from the experimental values to 18oC. n states the number of moles
represented for the portion of the equation that is being discussed. Cp stands for the
specific heat. This equation was calculated for ammonium nitrate, water, and the solution
of the two. The varience with calculating the energy change for the solution was that the
a combination of the moles and specific heats of the two portions of the mixture had to be
made as follows:
 Hsolution  [(nA N Cp, A N )  (nwater Cp,water )]   T
Lastly, knowing that the initial experiment did not have any heat or work added to the
reaction, a heat balance could be made to determine the heats of solution for the various
compositions of the ammonium nitrate and water.
(4)  H A N   Hwater   Hs ,18 C   Hsolution  0
All energy changes in the above equation are known except for the heat of solutions at
18oC. By manipulating the equation, calculations could be made to determine the heats
of solution as follows:
 H s ,18 C    H A N   H water   H solution
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EQUIPMENT
For this particular experiment, there were a few pieces of equipment and materials that
were needed for analysis. First, ammonium nitrate was needed to add into solution with
the water. A 5 mL vial was used to measure the amount of ammonium nitrate that was
being added to the various amount of water. There were four different thermometers
used to determine the temperatures for this experiment. The reason for using four
thermometers was to get an average value of temperatures for each measurement taken.
Also, needed for this experiment was a Styrofoam cup for the mixing of the ammonium
nitrate and water solutions. A 100 mL graduated cylinder was used to measure amounts
of water for the solutions.
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PROCEDURE
The procedure for this experiment was as follows:
1. Measure out 240mL of water and add to Styrofoam cup
2. Take temperature of water and record
3. Measure out one 5mL vial of ammonium nitrate (8.65g) and add to Styrofoam cup
4. Stir solution until ammonium nitrate is completely dissolved and record
temperature of solution
5. Clean out cup and begin again by decreasing water by 20mL
6. Repeat steps 2-5 until reaching 20mL of water.
After recording these measurements, the change in temperature can be determined, and
the various heats of solution can be calculated.
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RESULTS
Upon taking the various measurements of temperatures for the water and the solutions,
the change in temperature calculations could be made along with their associated errors.
These values are listed in the table that follows.
Table 1. Measured Data from Heat of Solutions Experiment
mL of Water
240.0
220.0
200.0
180.0
160.0
140.0
120.0
100.0
80.0
60.0
40.0
20.0
Average Temp.
of Water
20.4
20.9
21.3
21.3
21.5
21.3
21.3
20.8
21.4
21.2
21.3
21.3
Average Temp.
H20 & NH4NO3
18.3
18.6
18.2
18.6
18.5
17.9
17.3
15.9
15.3
13.6
10.1
2.5
Average T
2.0
2.3
3.1
2.7
3.1
3.4
4.0
4.8
6.1
7.6
11.2
18.8
Student t
0.44
0.55
0.55
0.55
0.66
0.77
0.77
0.623333
0.44
1.95E-15
0.11
0.22
A graphical representation of the overall temperature changes upon the addition of the
ammonium nitrate to the water is shown in the following figure. The graph also shows
the student t error bars that represent 99% accuracy.
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Temperature Change Upon Addition of Ammonium Nitrate
Temperature Drop (oC)
20.0
15.0
10.0
5.0
0.0
0.0
50.0
100.0
150.0
200.0
250.0
Water (m L)
Figure 1. Temperature Difference after Addition of Ammonium Nitrate
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The moles of ammonium nitrate and water for each solution had to be calculated. The
ammonium nitrate remained the same for each measurement taken (0.105 moles of
ammonium nitrate). The moles of water varied for each measurement. The table of those
values is as follows:
Table 2. Moles of H20 and NH4NO3
mL of Water
240.0
220.0
200.0
180.0
160.0
140.0
120.0
100.0
80.0
60.0
40.0
20.0
moles of H20
13.33
12.22
11.11
10.00
8.89
7.78
6.67
5.56
4.44
3.33
2.22
1.11
moles NH4NO3
0.105
0.105
0.105
0.105
0.105
0.105
0.105
0.105
0.105
0.105
0.105
0.105
Upon computing the moles of water and ammonium nitrate, the final calculations for the
energy changes for the various reactions taking place could be completed. This led to the
calculations of the heat of solutions for the various solutions at 18oC.
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Table 3. Heats of Solution Calculations
HA-N (J/mol
solute)
-7022.323
-7022.323
-7022.323
-7022.323
-7022.323
-7022.323
-7022.323
-7022.323
-7022.323
-7022.323
-7022.323
-7022.323
Hwater (J/mol
solute)
-130.973
-146.881
-150.944
-136.895
-130.044
-108.099
-90.567
-64.635
-63.164
-43.890
-30.653
-15.211
Hsolution (J/mol
solute)
474.564
800.554
253.909
795.446
648.729
-107.773
-1038.440
-2938.655
-3877.791
-6311.871
-11133.309
-21805.963
mL of Water
240
220
200
180
160
140
120
100
80
60
40
20
Hs @18oC
(kJ/mol solute)
6.679
6.369
6.919
6.364
6.504
7.238
8.151
10.026
10.963
13.378
18.186
28.843

HA-N represents the heat required for ammonium nitrate at 23oC (room temp.) to change
to 18oC. Hwater represents the amount of energy need for the temperature of the water to
change to 18oC. The Hsolution represents the energy required to force the solution to
change to 18oC.
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A graphical representation of the heats of solutions at 18oC can be seen in the figure that
follows.
Heats of Solution @18 oC
(KJ/mol NH4NO3)
Heats of Solution @ 18 oC
30.0
25.0
20.0
15.0
10.0
5.0
0.0
0
50
100
150
200
250
Water (mL)
Figure 2. Heats of Solutions at 18oC
The calculated values show that at dilution (complete saturation) the heat of solution for
ammonium nitrate in water was 29 kJ/mol A-N at 18oC. Perry’s Chemical Engineering
Handbook gives the heat of solution for ammonium nitrate at 18oC to be 27 kJ/mol A-N
at dilution.
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DISCUSSION OF RESULTS
The results show that as the amount of water increased in the ammonium nitrate
experiment, the temperature drop decreased as well. The results show that between
160mL and 240mL of water in solution, there was an exothermic reaction, which means
that the solution would have to give off energy to reach the temperature of 18oC. For the
rest of the reactions, there were endothermic reactions that would have to take place in
the solutions in order to reach 18oC. Figure 2 on the previous page shows that near
dilution the heat of solution for ammonium nitrate in water is approaching 30 kJ/mol A-N
at 18oC. With Perry’s Chemical Engineering Handbook being 27 kJ/mol A-N for the
heat of solution at 18oC, a %error was calculated to be 7%.
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CONCLUSIONS
Knowing that, from literature, the heat of solution for the experiment ran near dilution
should be 27 kJ/mol solute, it can be concluded that the experiment ran was successful in
reproducing the dilution data within 7%. Teams II generated a heat of solution of 29
kJ/mol solute near dilution as well as determined various heats of solution at other mol
fractions of the ammonium nitrate and water solution. For the purpose of this
experiment, it is safe to say that the data calculated can be used to determine heats of
solution for ammonium nitrate in solution with water.
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REFERENCES
Unit Operations of Chemical Engineering, 6th Ed., Warren L. McCabe, Julian C. Smith and
Peter Harriott, McGraw-Hill, Boston, 2001.
Perry’s Chemical Engineering Handbook. 5th Ed.. Robert H. Perry, Don W. Green and James O. Maloney
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