How Minnesotans Clear the Streets in Winter:
An Exercise in Freezing Point Depression
In section 11.5 of Zumdahl, you learn about the effect of adding a solute to a liquid: the
boiling point is elevated and the freezing point is decreased. In this experiment, you will
look at one of the most common applications of this idea: the use of chemicals to melt ice
on roads and highways. By exploring the ice-melting properties of different substances,
you may even come up with a better chemical for the job!
Introduction
City, county, and state governments have responsibility for clearing streets of ice and
snow every winter. Besides physically clearing the streets with snowplows, crews reduce
the risk of skids by dumping sand (for traction) and rock salt on the streets. The rock salt
acts to remove ice by lowering its melting point to the ambient temperature (or below).
But, as many Minnesotans have experienced, the salt can greatly accelerate the rusting of
iron and steel in cars and of the steel in steel-reinforced concrete. Salt is also detrimental to
concrete surfaces and it can kill roadside plants. A less-corrosive substitute for rock salt
would save taxpayers lots of money -- both for repair of highways and in rust damage to
their cars. In this experiment, teams of students will measure the effectiveness of rock salt
(NaCl) and some other chemicals to determine their effectiveness as deicing materials, and
evaluate their cost and environmental impact. Finally, your class may be asked to prepare
a report in the form of a poster summarizing the pros and cons of rock salt and the
possible alternative deicers studied by your lab section.
Dissolving a solute in a liquid both raises its boiling point and lowers its freezing point.
For ideal solutions of non-ionic solutes, the freezing point depression is calculated by
∆T = Kfm
where ∆T is the lowering of the freezing point, Kf is a constant that must determined
experimentally for each substance, and m is the molality of the solution (moles solute per
kg of solvent).
10-1
(∆T, as it is defined here and in Zumdahl, is a positive number.) The Kf constant for water
is rather small: 1.86 °C per molality unit. Other substances can have quite large Kf
constants: for example, the Kf of CCl4 is 30. °C•kg/mol. Since the freezing point of pure
water is 0 °C (or 32 °F), addition of any solute will lower the freezing point of pure H2O by
∆T, to a lower freezing point of Tf = 0 - ∆T. The new Tf will thus be below 0°C.
For ionic or ionizing solutes (for example HCl), the actual concentration of particles
depends on the molality of the solute (moles ionic compound per kg water) and the
number of ions in the formula unit. For example, a solution containing 0.01 moles MgCl2
per kg of water would have a “particle” or “colligative” molality of 0.03, because there
would be 0.01 moles Mg2+ ions and 0.02 moles of Cl– ions in every kg of solvent. An
additional complication arises because of ion-pairing effects, as described in Section 11.7 of
Zumdahl. The experimentally measured colligative molality is usually somewhat less than
expected because a small percentage of the ions in solution are believed to cluster
together, for instance as (MgCl)+ ions. Thus the concentration of ions in 0.01 molal MgCl2
may actually be less than 0.03 m. The freezing point (solid and liquid in equilibrium) of an
aqueous salt solution is thus
∆Tf = imKf
where m is still the molality of the solution and “i” is the number of moles of ions put into
solution per mole of salt dissolved. For MgCl2 , i ≅ 3.
Finally, the decrease in freezing point is limited by the solubility of the deicer chemical
in water at very low temperatures. The lowest melting point that can be reached using a
given solute is called the eutectic (from the Greek for “easily melted”).
Note: It is the endothermic melting of ice that is responsible for cooling the solution. For
water, ∆Hfus = +6.02 kJ/mol and was determined in Experiment 5.
How Does a Salt Dissolve in Ice and Cause it to Melt?
It is undoubtedly helped by a phenomenon called “surface melting.” Surface melting is
the formation of a very thin layer of liquid at the surface of a solid at temperatures below
the melting point of the solid. In the 1840s the great English scientist Michael Faraday
10-2
proposed that there is a microscopic layer of liquid water on the surface of ice and that it
persists well below the freezing point of water. Recent research confirms that “virtually all
solids undergo surface melting.”("Melting Below Zero,” by S. Wettlaufer and J. G. Dash,
Scientific American, Feb. 2000, p. 50-53.) It is reasonable to assume that salt will begin to
dissolve by forming a solution in this thin layer of water. Fortunately it is easier to observe
the ice-melting process and to predict the minimum temperature that can be reached as
more ice melts than it is to know precisely what is happening at the molecular level! We
can, however, trust that the "system" under study is moving towards its
thermodynamically most stable state.
Experiment Overview
In this Experiment, several properties of salts that are related to their suitability as deicers
will be studied. The data collected and the observations made for a given salt will be made
by a group of about four students and the results will be shared with all members of your
lab section. You will then compare and rank these salts as potential deicers.
Part I: Freezing Point Depression
Excess ice is added to a salt solution of known composition by mass. Ice melts, diluting the
solution. The solution will quickly reach a constant minimum temperature that is its
freezing point. The computer will record the temperature change on a temperature (yaxis) versus time plot. The minimum temperature reached is easily and accurately
measured. After measuring the freezing point of solutions of different concentrations,
make a plot of ∆T (a positive number!) vs molality. The slope of the line will be iKf.
Part II: The corrosive Properties of Salt Solutions
You will compare the corrosion caused by solutions of each deicer on an iron nail; on
concrete surfaces; and on an iron rod imbedded in concrete (the latter two were subjected
to 10 freeze-thaw cycles simulating the freeze-thaw conditions of a Minnesota winter).
Part III: Enthalpy of Solution
Finally, using the same techniques described in Experiment 5, you will measure the
enthalpy of solution, ∆Hsoln, of your salt (see Zumdahl, Section 11.2 for the theory behind
this measurement). What are the implications for a deicer’s effectiveness if it dissolves
exothermically versus endothermically?
The last steps are to carry out all indicated calculations, to collect, in tabular form, the
results for each deicer studied by your lab section, and to compare and rank these
potential deicers from most to least effective.
10-3
Week One: Gathering Data
Your TA will divide your lab section into five groups of 3 to 5 students per group.
He or she will then assign one of the following 5 deicers to each group: NaCl, KCl,
KC2H3O2, MgCl2, or CaCl2. If your group consists of 4 or 5 students, it will be most time
efficient to split up and use two computers. You’ll need to calibrate a temperature sensor
(Part IA) for each computer. One group should begin by preparing the two beakers for
Part IIA, then calibrating the temperature sensor and finally proceeding with Part III.
The second group will also do Part I A and will then proceed with trial 1 of Part IB.
All of Part IB needs to be done using the same computer, but each group member should
be part of at least two freezing point depression trials—first as an “assistant/learner” and
then as the “leader/trainer”. All group members will need to keep a record of all data
collected on Data Sheet I. If you calculate and plot ∆T vs. m as you go, and if you obtain a
good straight line for trials 1 through 4, you may omit trial 5 and proceed to trial zero.
Omitting trial 5 will significantly reduce the amount of chemical used.
Week Two: Comparing Data on Deicers and Evaluating their Effectiveness
1. Compare the damage to nails in each deicer after one week.
2. If not completed during week one, compare concrete samples.
3. Each group in your lab section will share its experimental results and observations
for Parts I, II and III. See Post-Lab Questions 2 through 4 for further directions.
EXPERIMENTAL SECTION
Waste Treatment
Because road deicing chemicals would be applied by the ton, we have chosen
environmentally safe substances to study. Unless announced to the contrary by your TA,
all solutions and dry solids may be rinsed down the drain with running water.
Part I Freezing Point Depression
A.
Calibration of Temperature Probe
You will be calibrating the temperature probe for your instrument using two known
temperatures – that of boiling H2O and that of freezing H2O. You will need to obtain one
or more temperature probes for your group.
10-4
1.
Using your Bunsen burner and an Erlenmeyer flask, heat deionized H2O to a
"gentle" boiling. In a beaker put a 50-50 mixture of cold deionized H2O and ice and
let the temperature equilibrate to 0º C.
2.
Plug your temperature probe into port one of the serial box interface. Open the
Logger Pro program, go to the Setup menu and choose the Sensors option. Click
on the “Port One” icon and then under the Sensor menu choose “Temperature –
Direct Connect” (near the bottom of the list). Now choose the Calibrate menu and
again click on the “Port One” icon. Click on the “Perform Now” button
3.
Place your probe into the boiling water such that the tip of the probe is just
submersed in the boiling water (you may secure the probe in place using an
additional clamp). Let the probe equilibrate in the boiling water for 5 minutes and
then, when the voltage reading is constant, enter the temperature of the boiling
water (100º C) for “Value One” and click on the “Keep” button.
4.
Next, place the temperature probe in the ice water and again let the probe
equilibrate for 5 minutes. You should stir the ice water to ensure there is no
localized warm spot near the probe. When the voltage reading has become
constant, enter the ice water temperature (0º C) for “Value Two” and click on the
“Keep” button.
5.
Your probe is now calibrated. Click on the Details menu to obtain the slope and
the intercept of the line that defines your calibration curve. Write these values on
Date Sheet I. Any voltage read by the temperature probe will now be assigned a
temperature based on the linear calibration curve just calculated.
B.
Determining the freezing point depression (Data Sheet I)
1.
Fill your squirt bottle with ice and deionized H2O so you have cold H2O available.
2.
Record the mass of the empty Styrofoam cup. Zero the balance and add to the cup
an appropriate amount of your deicer and record its mass. Try to get within ±10%
of the recommended sample size. Again zero the balance and add approximately
50 mL of cold deionized H2O and record the mass of the added H2O. Stir the
mixture until the deicer is completely dissolved.
10-5
3.
Place the Styrofoam cup with the solution into an ice water bath and place the
temperature probe in the deicer solution making sure that the probe does not
contact the sides or bottom of the cup.
4.
Go to the Set Up menu and choose Data Collection. From the Sampling menu
adjust the run to 2 minutes with a sampling rate of 60 readings per minute. Adjust
the y axis to display the range 5 ºC to –20 ºC (clicking on the y axis will bring up a
menu allowing you to make this adjustment).
5.
Click on the Collect button to begin the run. After you start the run, immediately
begin putting ice into the solution (enough ice so that when melting is complete
there is still some ice remaining in the cup). You should constantly stir the solution
during the run. When the run has ended, the temperature should have reached a
constant minimum
6.
Record the minimum temperature reached during the run.
7.
Immediately weigh the cup with the ice and deicer solution. Quickly decant any
liquid from the cup, being careful not to lose any ice, but also making sure to
remove all liquid. Again weigh your cup with the remaining ice. You can obtain the
mass of the ice added to the deicer solution and the mass of the ice melted by
subtraction.
8.
Group members not involved in data collection should calculate and plot ∆T vs.
molality. After all the points are plotted, repeat any trials that appear to be in error.
9.
Be sure to thoroughly rinse out the styrofoam cup and dry it after each run. After
your group has completed all the runs with the deicer, with the possible exception
of Trial 5, repeat the procedure one more time using no deicer and record the
numbers under Trial 0. The minimum temperature reached for Trial 0 should be
0°C ± 0.5°C.
10.
Complete Post-Lab Question 1.
10-6
Approximate Sample Sizes for Deicers
Deicer
Trial 1
Trial 2
Trial 3
Trial 4
Trial 5
NaCl
2g
4g
6g
8g
10 g
KCl
1g
2g
4g
6g
7.5 g
KCH3CO2
1g
2g
3g
5g
7g
MgCl2
0.5 g
2g
3g
4g
5g
CaCl2
1g
2g
3g
4g
5g
Part II: Corrosive Effects of Deicers
A.
The Corrosive Effects of Deicers on Iron Nails
Only one person in your group needs to carry out the following directions and to do it
only for the group’s deicer. Therefore, during week two, all of the groups' beakers need to
be placed together so everyone in your lab section can examine all the nails and record
their observations.
1.
Take two iron nails and examine them carefully for signs of rust. If any rust is
present, clean it off with sandpaper or steel wool. The nails don’t have to be shiny,
just free of orange or brown deposits.
2.
In a small, labeled beaker, prepare a solution containing 2.0 g of deicer (this may be
weighed on a triple beam balance) in 20 mL of water. Place one nail in this solution
and one in 20 mL of DI water. Try to lean each nail against the side of the beaker so
part of it sticks out into the air.
3.
After approximately 30 minutes, after one hour, and again at the end of the lab
period, examine the nails carefully. Is either of them showing signs of rust? Report
your observations on Data Sheet IIA.
4.
Leave the beakers in your drawer covered with a watch glass or with a larger,
inverted beaker. It’s okay if all the water evaporates during the week. The nails
will have been exposed to the liquids for at least a few days. Next week place your
labeled beakers on a counter with the other nail samples from your lab section and
record your comparisons.
10-7
B.
The Corrosive Effects of Deicers on Concrete and Iron Surfaces During Repeated
Freeze - Thaw Cycles
Concrete scaling refers to damage done to concrete surfaces under various conditions. The
amount of scaling is rated from 1 to 5, with 1 being no surface damage during the study
and 5 being the most significant damage.
Two sets of concrete samples were prepared in plastic containers. One set consists of
concrete in 1300 mL beakers. In the second set, the concrete samples were placed in 800
mL containers and a U – shaped iron bar was imbedded in the concrete. Each concrete
sample was covered with one of 9 different solutions and was then subjected to 10 freezethaw cycles between about -17°C and room temperature. Because of the time required to
prepare the concrete samples and to put them through a number of freeze – thaw cycles,
the samples were prepared and treated outside of the regular laboratory and over a
period of several weeks. The resulting concrete samples will be on display during the two
weeks spent on Experiment 10.
1.
Examine the concrete samples which have no iron bars present, estimate the
amount of scaling on each sample and report the results, which will be subjective,
on Data Sheet IIB.
Do not touch the concrete surfaces!
2.
Also describe the extent of corrosion on the U – shaped bars on a scale of 1 to 5 and
report the results on Data Sheet IIB. Once again the results, while reported to the
best of your ability, are qualitative. For a given deicer, did the presence of the iron
bar affect the amount of scaling on the concrete surface?
3.
Look around campus and around town for places where rusty iron bars and flaked
concrete reveal damage to bridges, parking ramps, roadways, etc. (Hint: A photo
of serious rust or crumbling concrete, with the location noted, would be a good
addition to your lab report or to a poster on deicers.
10-8
Part III: Determining the Enthalpy of Solution, ∆Hsoln
Many factors go into making a deicer both effective and practical. An obvious one,
perhaps so obvious we usually don’t mention it, is solubility. Since freezing point
depression is proportional to the molality of the dissolved solute, a solute with low
solubility in water would be useless as a deicer. Therefore, you wouldn’t expect to see
CaCO3 used as a deicer because of its low solubility although it has other problems as
well.
Another factor to consider in selecting a deicer is its enthalpy of solution. If a solid
dissolves exothermically, it will provide heat to its surroundings and tend to increase the
amount of ice melted per gram. From Section 11.2 of your text (Zumdahl, 5th Ed.), we see
that the dissolving process can be considered to take place in three steps. According to
Hess’ Law, the overall enthalpy of solution is the sum of the enthalpies of the three
individual steps. The first two steps require the separating of solvent particles and solute
particles (or “expanding” the solvent and solute) and are always endothermic. The third
step is the coming together of the solute and solvent particles. If this step is sufficiently
exothermic, the overall process is exothermic. (Consider the rule learned earlier that “like
dissolves like”. The stronger the attraction between solvent and solute particles, the more
exothermic, or less endothermic, we would expect the third step to be.)
For an ionic solid such as NH4Cl, the overall reaction for dissolution is
NH4Cl(s) + H2O(l) → NH4+(aq) + Cl-(aq)
?
∆Hsoln
If a salt dissolves endothermically, will the temperature of the resulting solution increase or
decrease? Will the ∆Hsoln you calculate be a positive or a negative number?
In this part of the experiment you will determine the approximate enthalpy of solution for
your deicer. During week two, you will also collect ∆Hsoln values for the other deicers and
use these values to help rank the deicers from the salt which dissolves with the largest
negative enthalpy of solution to the one which dissolves with the largest positive enthalpy
of solution.
10-9
Part A. Determination of the Calorimeter Constant
Refer to Experiment 5 for a sample calculation of a calorimeter constant. Your
calorimeter is not perfectly insulating; that is, a small amount of heat is either absorbed or
released by the calorimeter each time you run an experiment. The amount of heat gained
or lost by the calorimeter per degree change in temperature is called the calorimeter
constant. It is the heat capacity of your calorimeter. You will determine this constant by
mixing warm water with cool water in your calorimeter.
1.
Using your small graduated cylinder, measure approximately 5 mL of cool water
(approximately 10°C – you may use the water from your ice bath used in the
calibration of your temperature sensor) and record its volume to the nearest 0.02
mL. Pour the water into a calorimeter well. Using a clamp place the temperature
probe into the well such that it is reading the temperature of the water but is not
touching the sides or bottom of the well.
2.
Go to the “set up run” menu and adjust the run to 5 minutes with a sampling rate of
60 samples per minute. Begin the run and then quickly, using the same small
graduated cylinder as before, measure approximately 5 mL of hot water (about 30 35°C) from the hot-water tap and record its volume to the nearest 0.02 mL.
Immediately before mixing the two solutions read and record the initial temperature
of the hot water using the digital thermometer.
3.
Pour the hot water into the well and stir until the temperature stops rising. Click on
the stop run button. To obtain the initial temperature of the cold water, use the
mouse to highlight the portion of the graph immediately prior to the addition of the
warm water and then click on the “stat” button (a box will appear that will give you
the average temperature of the water for the portion of the graph you highlighted).
Repeat this procedure to obtain the maximum temperature reached. This is the final
temperature of the solution. Record these initial and final temperatures in your
notebook. For all temperature measurements, do not rest the probe on the bottom
of the flask!
10-10
Part B. Determination of ∆Hsoln for your Deicer
1.
Refer to Table 10.1 to confirm that the suggested mass of sample will dissolve
completely in 10 mL of water.
2.
Use a small graduated cylinder to measure 10 mL of water into a calorimeter well.
Record the exact volume (to the nearest 0.02 ml) on Data Sheet III. Suspend a
calibrated temperature sensor in the water.
Suggested Sample Sizes:
3.
CaCl 2 0.6 to 1.0 g
KCl
1.5 to 2.0 g
NaCl 1.25 to 1.75 g
KC2H 3O2 1.5 to 2.0 g
MgCl2 0.25 to 0.3 g
Weigh the indicated amount of salt in a small, clean, dry beaker and record its mass.
Begin the run and when the temperature sensor reading has reached a constant
temperature add the salt to the well. Stir with the temperature sensor until the
temperature reaches a maximum or a minimum. Record the initial temperature of
the water as well as the maximum or minimum temperature reached. You can now
predict the sign of ∆Hsoln
4.
Rinse the well thoroughly with DI water and repeat the procedure, adjusting the
mass of sample used, sampling rate, etc. if necessary.
5.
Referring to the sample calculation below, calculate the enthalpy (heat) of solution for
your deicer in J/g and kJ/mol. If the specific heat of your deicer is not given in Table 10-1,
use the specific heat of water, 4.184 J/g°C, for your solution of dissolved deicer.
Table 10.1 Solubility (g/100mL H2O) and Specific Heat (J/g°C) of an Aqueous Solution of
Selected Ionic Compounds
Compound
Solubility
Solubility
Specific Heat (J/g°C )
(°C )
(100°C )
of Salt Solution
NH 4 C l
29.7 (0°C)
75.8
3.31
CaCl 2
74.5 (20°C)
159
3.75
KCl
34.7 (20°C)
56.7
3.45
MgCl 2
54.2 (20°C)
72.7
3.94
NaC2H 3O 2
119 (0°C)
170
3.32
NaC 2H 3O 2•3H 2O
76.2 (0°C)
138 (50°C)
3.32
KC2H3O2
253 (20°C)
492 (62°C)
-----------
NaCl
35.7 (0°C)
39.1
3.60
10-11
Sample Calculation: Determine the Enthalpy of Solution of Ammonium Chloride.
A student is asked to add ~2g of NH4Cl to 10mL of water and to use the change in
temperature to calculate ∆Hsoln . The reaction below corresponds to the enthalpy of
solution of ammonium chloride.
NH 4Cl(s) + H2O(l) → NH4+(aq) + Cl-(aq)
∆H soln
A student first checked the solubility of NH4Cl to make sure that a 2 g sample would
dissolve in 10 mL of water. She reasoned that because 2.97 g will dissolve in 10 mL of
water at 0°C and since the solubility of NH4Cl increases with temperature, a 2 g sample
should readily dissolve at room temperature. She then stirred 2.105 g of NH4Cl into 10.00
mL of DI water and observed that the temperature changed from 24.2°C to 11.1°C. The
calorimetry constant for the well used was 5.8 J/°C. If qsoln is the heat transfer associated
with the dissolving of the salt sample and q is the heat lost or gained by the solution of
water and salt, then
qsoln + q + qcal = 0
qsoln + (mass of salt + mass of water)(Csp solution)(∆T) + (5.8 J/°C)(∆T) = 0
qsoln + (2.105g NH4Cl + 10.0 mLx 0.997g/mL)(3.31 J/g°C)(11.1°C – 24.2°C) +
(5.8 J/°C)(11.1°C – 24.2°C) = 0
qsoln − 524 J
− 76 J = 0
and
qsoln ≅ 6.00 x 102 J
∆H(J/g) = (6.00 X102 J)÷(2.105 g) = 285 J/g
∆H (kJ/mol) = (285 J/g) (1 kJ/ 1000 J) (53.49g/mol) = 15.2 kJ/mol
10-12
Name:
_________________________________________
Date:___________________
Pre-Lab Questions for Experiment 10
Four students have completed Part I of Experiment 10. Their deicer was NH4Cl (molar
mass = 53.49 g/mol). The data for trials 1,2, and 3 is given in the following table.
For trial 0, the freezing point of water, the temperature was 0°C and m = 0. Refer to this
data to answer the questions below.
1. Complete the Table below. Show your work for parts c, d, e, and f.
Mass of deicer used
Mass of water added
Mass of ice melted
Minimum temperature reached
a) Mass of water in final solution
b) kg of water in final solution
c) Moles of deicer used
d) ∆T
e) Molality of final solution*
f) Grams of ice melted/ g deicer
* The value required here is "m" not "i x m"
c)
d)
e)
f)
10-13
Trial 1
Trial 2
Trial 3
1.920 g
4.119g
5.861g
50.525g
49.872g
48.955g
10.584g
23.916g
32.074g
-2.11°
-3.88°C
-4.70°C
2.
Prepare a plot of ∆T vs. m. Include trials 0 through 3. Use Excel or the graph paper
at the back of the Manual. Each part (a, b, and c below) must be entirely correct to
receive credit for that part.
a)
Put your name on the graph, provide a title, and label the axes.
b)
Plot the 4 data points and draw the “best” straight line.
c)
Calculate the slope of your line, ∆T = imKf
d)
Calculate “i” from this slope. (Slope = iKf)
e)
What is the theoretical value of "i" for NH4Cl? How does your value compare
to the theoretical value of “i” for NH4Cl?
f)
Use the experimental value of “i” to determine the “colligative molality, i x m,
where m is the calculated molality of the compound (not of the ions.)
10-14
Name:
_________________________________________
Partner(s): _________________________________________
Date:___________________
TA: ____________________
Data Sheet I: Data and Calculations for Part I
Calibration Data:
Slope: ______________
Intercept: ______________
Deicer: ______________
Molar Mass of Deicer: ___________
For each trial, record the experimental data.
Trial
0
1
1. Mass of cup
~2 g
2. Mass of deicer
0
3. Mass of cold H2O
~ 50 g
2
3
4
5
3
4
5
4. Minimum temperature
5. Total mass after adding ice
~65 g
xxx
6. Mass of cup + remaining ice
For each trial you carried out, complete the table below:
Trial
0
1
2
7. Mass of deicer (2)
xxx
8. Mass of cup, cold water,
and deicer (1+2+3)
9. Mass of ice added (5-8)
xxx
10. Mass of ice remaining
(6-1)
11. Mass of ice melted
(9-10)
12. Mass of final solution
(5-6)
13. Mass of water in final
solution (12-2)
xxx
xxx
xxx
xxx
xxx
Complete the table below (see Post-Lab Question 1.) Show complete calculations for 2 trials
on the next page.
Trial
0
Moles of deicer used
xxx
Kilogram of water in the final
solution
Molality of final solution
xxx
∆T = |0°C - Tf |
Grams of ice melted per gram
of deicer
1
2
xxx
xxx
xxx
10-15
3
4
5
Show the calculations required for any two of Trials 1 through 4.
Trial Number: _______
Trial Number: _______
10-16
Name:
_________________________________________
Partner(s): _________________________________________
Date:___________________
TA: ____________________
Data Sheet IIA: Observation of the Corrosive Effects of Deicers on
Iron Nails
Write a brief description of the nail surfaces
Time Exposed
Plain H2O
H2O + Deicer
After 30 minutes
After 1 hour
End of Lab (total time nails
exposed = _______)
After 1 week
Corrosion Results After One Week for the Remaining Deicers
See Data Sheet IIB for directions for "rating" the samples.
Deicer Sample Observation
10-17
Rating
Data Sheet IIB:The Corrosive Effects of Deicers on Concrete and the
Surfaces of Iron Embedded in Concrete
CAUTION: Do not physically disturb the surface of the iron rod or the surface of the
concrete during your examination of the samples.
Concrete Sample Rating
Concrete
Rating
1
no scaling,
surface intact
2
Some
surface
“powder”
3
Scaling
penetrates
the surface
4
Fairly deep
corrosion,
large cracks
and flakes
5
Soft
surface,
deep
scaling
Iron Bar Sample Rating
Iron Bar
Rating
1
Entirely
intact,
no color
change
2
Orange
surface
splotches
3
Mostly
orange, size
intact
4
Entire bar
orange,
some
flaking
5
Bar visibly
smaller as
pieces flake
off
Use your best judgement in assigning the ratings.
Deicer Solution
Concrete Sample Rating
Water
10-18
Iron Bar Sample Rating
Name ___________________________________________ Date ____________________
Partner __________________________________________ TA _____________________
Experiment 5 DATA SHEET
Part A. Calibration of the Temperature Probe
1.
a) Record the trial Calibration Curve Equation here: ______________________________
b) Record the Calibration Curve Equation here: _________________________________
2.
Does the probe measure temperature directly? Explain.
Part B. Calorimeter Constant Determination
Trial 1
Trial 2
1. Volume of cold water
________ mL
________ mL
2. Mass of cold water
________ g
________ g
3. Temperature of cold water
________ °C
________ °C
4. Volume of hot water
________ mL
________ mL
5. Mass of hot water
________ g
________ g
6. Temperature of hot water
________ °C
________ °C
7. Final temperature after mixing (Tf)
________ °C
________ °C
8. Calorimeter constant
________ J C° −1
________J C° −1
____________ J C° −1
9. Mean calorimeter constant
Show complete calculations of the calorimeter constant here.
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Data Sheet III: Determination of the Enthalpy of Solution of a Salt
Deicer: Formula _______________________ Molar Mass ________________________
Estimated grams of compound that will dissolve in 10 mL of water at 25°C: _____________g
Trial 1
Trial 2
1.
Volume of H2O used
________ mL
________ mL
2.
Mass of water used
________ g
________ g
3.
Mass of sample used
________ g
________ g
4.
Initial temperature of water
________ °C
________ °C
5.
Final temperature of solution
________ °C
________ °C
6.
q soln
________ J
________ J
7.
Enthalpy of solution per gram
________ J/g
________ J/g
8.
Enthalpy of solution in kJ/mol
________ kJ/mol
________ kJ/mol
9. Average value of the enthalpy of solution
________ kJ/mol
Show a complete calculation of the enthalpy of solution below:
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Name ___________________________________________ Date ____________________
Partner __________________________________________ TA _____________________
Post-Laboratory Questions for Experiment 10
1.
Data Treatment for Part 1
a)
b)
c)
d)
Carry out the necessary calculations and complete the last table on Data Sheet I.
Prepare a graph for plotting ∆T vs. molality, m for trials 0 through 4 or 5. Select
a scale such that the graph fills over half of a page. Give the graph a title and
label the axes. (Whether the graph is done by hand or using the computer, any
missing information will result in a loss of all points for b).
Graph your 5 or 6 data points, including DI water, and draw the best straight
line through the data points. The formula for your straight line is ∆T = imKf .
(An error in graphing your data points or in drawing the best straight line will
result in the loss of all points for c).
Determine the slope of ( i x Kf ) and calculate “i”. Compare this to the
theoretical value of “i”. For example, we expect i ≤ 3 for MgCl2. Show your
work.
2. Summary of Results
Use the computer to prepare a table that summarizes the experimental results and any
other information available to your lab section on the 5 deicers. For each deicer,
include as much of the following information as possible. “Average values” may refer
to the average of several trials, as is the case for ∆Hsoln, or, as in the case of rating
surface damage, take the average of the ratings assigned by each of the groups. You
need to submit only one table per group.
a)
b)
c)
d)
e)
f)
Slope of the plot of ∆T vs molality.
Value of “i”
Average mass of ice melted per gram of deicer
Solubility at 0 or 20°C (g/100 g solvent)
Average ∆Hsoln in kJ/mol
Average rating of the surface damage (corrosion or scaling) of the concrete
surfaces, iron bars, and iron nails
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3.
Evaluations and Conclusions
What do you think is the “best” deicer for Twin Cities streets and highways? Why?
Base your answer on the information obtained in this Experiment and summarized in the
table in Question 2. Explain or justify your choice. If there is a consensus within your
group, the group may submit a single table and paragraph here. However, if your opinion
differs from the group’s, you may decide to write a “dissenting” opinion. Write the name
of each contributor on all joint work.
4.
Other Important considerations
Your Table covers only part of the information that needs to be considered when selecting
a deicer. List at least 3 other factors that should be a part of the decision.
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Name:
_________________________________________
Partner(s): _________________________________________
Part V:
Date:___________________
TA: ____________________
Group Poster Project on “Selecting the Best Deicer for
Minnesota Roads”
(100 points: All contributing group members receive the same grade.)
Prepare a poster using poster board of standard size and thickness. (There will be a 10
point deduction from your assigned grade for failure to use the designated poster board.)
Poster board is available from the lab stockroom for about $1.00 per sheet. The poster
should be understandable to someone who has studied Chapter 11 in Zumdahl.
Use your answers to the post-lab questions, the questions below, and any other
information you have obtained to prepare your group’s poster. Your TA may request an
oral report from each group in addition to or instead of the preparation of a poster.
1.
Look up the cost of each deicer in a chemical catalogue. Use the cheapest possible
price, which will probably be for the largest quantity and the lowest purity. The
prices paid for very large quantities (thousands of tons), as would be needed by a
municipal or state highway department, would be considerably lower, but catalog
prices will hopefully at least give a relative ranking. You might also look for prices
on the Internet or in a chemical trade magazine. Cargill, a privately held Minnesota
based company, is a major supplier of deicers and does extensive research on them.
(Companies often do not quote prices to the public but negotiate them directly with
customers such as MNDOT who use very large amounts of a product.)
For the poster, search the library and Internet for information about the
environmental effects of deicer chemicals. You might visit the Web pages of various
highway departments and also nature magazines. Reference any web materials by
reporting both the complete address and the date you accessed the material.
a)
Which deicers are the least expensive in terms of cost per ton and cost per
pound?
b) Compare the deicers as mass of ice melted per dollar
2.
Economists sometimes talk about “hidden costs” - costs that do not show up
immediately, but that do ultimately add to the “price” of a deicer. For snow
removal, some of the obvious expenses (besides buying the deicer) are the payroll
for snowplow and truck drivers, fuel and maintenance for the trucks, etc.
However, most of these costs will be similar for each deicer compound.
Environmental damage during mining or production, pollution and any of the
resulting environmental degradation are examples of hidden costs.
3.
Select the “best” deicer from those studied in this experiment and write a brief
paragraph to support and explain your choice.
4.
Work in groups of 4 to 6 to prepare a poster which will name the “best” deicer for
Minnesota roads and which uses all available information to support your decision.
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