# Alison Poulson DETERMINING THE MOST EFFICIENT ICE

advertisement ```Alison Poulson
DETERMINING THE MOST EFFICIENT ICE-MELT
Introduction
In much of North America, winter snowfall causes accidents and injuries. To minimize
these problems, cities shovel roads and use ice-melt to keep them clear. Ice-melt is usually a salt
and lowers the melting point of H2O by changing the concentration of ions suspended in it. Any
water-soluble substance will lower the melting point of ice, but for ice-melt to be effective the
depressed melting point must be lower than the air temperature. There are many factors involved
in choosing a commercial ice-melt, but a key question is whether dissolving the substance is an
endothermic or an exothermic process. A more exothermic process releases enough heat to
continue melting the ice around it. Dissolving requires energy to break apart the lattice structure
of the solid salt, and then gives off energy as new bonds are formed with the water. Dissolving
can be exothermic if stronger bonds are formed than are broken, or endothermic if stronger
bonds are broken than are formed. In this experiment, it was predicted that anhydrous salts
would have the most exothermic reactions with water and thus be the best ice-melts, because
they do not have any bonds with water in the solid form to break in the dissolving process.
Experimental
This lab used a Styrofoam cup calorimeter (with a heat capacity of 55.8) and the
Labworks program to obtain temperature data. The mass of the cup was measured first. 40mL
of water was obtained in a graduated cylinder and added to the cup, and the cup and contents
were weighed. The water temperature was measured for three minutes. In the meantime,
approximately 2.0g of solute was set aside. After three minutes, the solute was added all at once
Alison Poulson
and the solution was stirred. The temperature was monitered for another 600 seconds. A graph
of the temperature vs time was used to obtain the initial and final temperatures of the solution,
and the solution was weighed to find the exact mass of the solute and of the solution. Using
these data, the enthalpy of each reaction was calculated.
Results
The following graph of the reaction of CaCl2 is a sample of the graphs of temperature vs
time obtained for each solute. In each case, a linear fit was made for the points before the time
of mixing, and the linear equation was solved for the time of mixing to give the initial
temperature. Another linear fit was made after the temperature steadied after mixing, and that
equation was solved for the time of mixing to give the final temperature. Only the change in
temperature is recorded for each solute in the table below.
25
y = -0.0015x + 23.468
R² = 0.9892
20
Series1
15
Series2
y = 0.002x + 16.166
R² = 0.9559
Series3
10
Linear (Series1)
Linear (Series2)
5
0
0
200
400
600
800
1000
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Solute
NaCl
NaC2H3O2
CaCl2*2H2O C6H12O6
CaCl2
NaC2H3O2*3H2O
0.0323
0.0222
0.0127
0.0102
0.0163
0.0125
41.24
41.28
41.58
41.08
39.98
41.06
182.6
182.5
182.6
203.0
182.6
182.5
1.8854
1.5343
-0.6101
6.6635
-1.5123
Moles of
Solute (mol)
Mass of
Solution (g)
Time of
Mixing (s)
Change in
Temperature -0.8854
(◦C)
The data above was used to calculate qrxn using the following formulas:
qrxn = -qsoln + -qcal
qsoln = (mass solution)(specific heat capacity water)(change in temperature)
qcal = (calorimeter constant)(change in temperature)
The enthalpy change for the reaction was found using the following formula:
ΔH = qrxn/molsolute/1000
The work is shown below for a sample, the reaction of CaCl2:
qrxn = -(39.98g)(4.184J/g◦C)(6.6635◦C) + -(55.8J/◦C)(6.6635◦C) = -1486J
ΔH = -1486J/0.0163mol/1000 = -91.2kJ/mol
qrxn (J)
NaCl
NaC2H3O2
CaCl2*2H2O C6H12O6
CaCl2
NaC2H3O2*3H2O
202.2
-431
-353.1
-1486
343.8
138.7
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ΔH
6.26
-19.4
-27.8
13.6
-91.2
27.5
(kJ/mol)
Discussion & Conclusions
The results of the calculations match the experimental data, as reactions that resulted in a
higher final temperature produced a negative enthalpy. The data seem to be accurate; there are
no surprising or outlying results. A high-magnitude negative number for ΔH indicates a highly
exothermic reaction. Therefore, according to these results, CaCl2 is the best ice melt. It releases
91.2kJ of energy for every mole used. As predicted, it releases less energy in its hydrated form –
only 27.8kJ/mol – but this is still more energy than any other ice melt releases.
NaC2H3O2 also follows the hypothesis. In its anhydrous form, it releases 19.4 kJ of
energy per mole. Its hydrated form, by contrast, requires 27.5kJ/mol to melt. This is the most
energy required by any of the solutes tested, so hydrated NaC2H3O2 is not appropriate for ice
melt.
Ice melt is often referred to as “salt,” NaCl, and blamed for damage to cars. However,
NaCl is not a very efficient ice melt because it too requires energy to melt. Glucose is even less
useful as an ice melt than NaCl. Overall, CaCl2 is the best choice for ice melt in terms of the
heat it produces. However, cost or environmental considerations may play a role in determining
which ice melt is finally chosen.
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