CHEM&153
WINTER 2010
SOLUTIONS
Adapted from Chemistry with Computers, Vernier Software
http://www.phas.ubc.ca/~year1lab/p100/LoggerPro/LoggerPictures/Toolbar.jpg
LEARNING OBJECTIVES: After completing this experiment, you should feel comfortable:
•
Using temperature probes and LoggerPro software to investigate the effects of temperature on
the solubility of a salt.
•
Analyzing a solubility graph to determine whether a solution is saturated or unsaturated.
•
Relating van’t Hoff factors for different ionic compounds to the freezing point of a solution.
•
Understanding the real behavior of ions in solution.
•
Measuring the freezing points of pure liquids and solutions.
•
Using a colligative property to determine an accurate molar mass of a compound.
TO EARN YOUR FINAL STAMP: The following items must be completed in lab. You may
complete the entire assignment in the lab; this reflects the minimum required to earn your final
stamp.
 Collect the data and acquire the graph for Part 1: Effect of Temperature on the Solubility of a
Salt and complete data table 1.
 Collect the data and acquire the graph for Part 2: Effect of the van’t Hoff Factor on the
Freezing Point of a Solution, and entering this in data tables 2 and 3. The data tables do not
need to be completed, but items 1, 3 and 6 should be entered into each table.
 Collect the data and acquire the graph for Part 3: Using Freezing-Point Depression to Find
Molar Mass and complete data table 4.
Turn in only the data tables and prelab on pages 10-15 and your graphs from
parts 1, 2 and 3.
Helpful info: Below is a key for the shortcut menu displayed by LoggerPro. Some of these buttons
will come in handy in this lab!
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Part 1: Effect of Temperature on the Solubility of a Salt
The solubility of a substance is dependent on a variety of criteria. For ionic compounds, recall that it
depends upon the charge on the ions and the distance between nucleii. In addition, temperature plays
a role in how much of a salt (ionic compound) will dissolve. Generally, as temperature increases, the
solubility increases. Measurements of solubility are typically done in g of solute per 100 g of H2O or
as Molar solubility in moles solute per liter of solution, Molarity.
In this experiment, you will study the effect of changing temperature on the amount of solute that will
dissolve in a given amount of water. Water solubility is an important physical property in chemistry,
and is often expressed as the mass of solute that dissolves in 100 g of water at a certain temperature.
In this experiment, you will completely dissolve different quantities of potassium nitrate, KNO3, in
the same volume of water at a higher temperature. As each solution cools, you will monitor
temperature using a computer-interfaced Temperature Probe and observe the precise instant that solid
crystals start to form. At this moment, the solution is saturated and contains the maximum amount of
solute at that temperature. Thus each data pair consists of a solubility value (g of solute per 100 g
H2O) and a corresponding temperature. A graph of the temperature-solubility data, known as a
solubility curve, will be plotted using the computer.
Figure 1
MATERIALS
Computer
Vernier computer interface
Logger Pro
Temperature Probe
2 utility clamps
four 20 X 150 mm-test tubes
test tube rack
ring stand
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hot plate
stirring rod
potassium nitrate, KNO3
distilled water
400-mL beaker
10-mL graduated cylinder or pipet
250-mL beaker
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Read all instructions before starting. Be careful with water and solutions next to electrical
connections and computers.
General Process
In Part 1 you will be finding the temperature at which the KNO3 begins crystallizing out of a
saturated solution. You will be dissolving different amounts the potassium nitrate in 5 mL of water so
that the concentrations are different. Each solution will have to be heated so that the solid completely
dissolves. Stirring is very important.
PROCEDURE
1. Obtain and wear goggles.
2. Label four test tubes 1-4. Into each of these test tubes, measure out the amounts of solid shown in
the second column below (amount per 5 mL). Note: The third column (amount per 100 g of H2O)
is proportional to your measured quantity, and is the amount you will enter for your graph.
Test tube number
Amount of KNO3 (in g)
used per 5 mL H2O
1
2
3
4
2.0
4.0
6.0
7.0
Amount of KNO3 used
per 100 g H2O
(use in computer spreadsheet)
40
80
120
140
3. Add exactly 5.0 mL of distilled water to each test tube. You may measure this volume or assume
a density of 1.0 g/mL for water so you may add this by mass.
4. Prepare the computer for data collection by opening the Vernier software program, Logger Pro
and then follow the pathway: File  Open Experiments  Chemistry with Computers  12
Temp and Solubility. The y-axis has solubility (in g/100g), rescale the axis from 0 to 200 g.
The x-axis should have temperature scaled from 0 to 100°C.
5. Do this on a lab bench, not next to the computers! Fill a 400-mL beaker three-fourths full of
tap water. Place it on a small hot plate situated on (or next to) the base of a ring stand on a lab
bench. This should not be near the computer!
Note: In order to dissolve all of the KNO3, gently heat test tubes 1 (about 30°C) and 2 (40-45°C), but
tubes 3 and 4 need to be heated to a higher temperature. While the water is warming up, you can heat
up test tubes 1 and 2. Use your glass stirring rod (or the temperature probe) to stir the mixture until
the KNO3 is completely dissolved. Do not leave the test tube in the water bath any longer than is
necessary to dissolve the solid. Be sure to clean the stirring rod between solutions.
Eventually, the water bath will need to be heated to about 90°C for test tubes 3 and 4. Adjust the heat
to maintain the water at this temperature. You will need this water bath for parts 2 and 3 of the lab.
Place a thermometer in the water bath to monitor the temperature. Place an unplugged probe in the
water bath to warm the probe.
CAUTION: To keep from damaging the Temperature Probe wire, unplug it from the LabPro unit
and hang it over another utility clamp pointing away from the hot plate, as shown in Figure 1.
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6. Lower one of the test tubes into the water as shown in Figure 1, to dissolve the KNO3. You may
use an unplugged temperature probe to stir the solution. Be sure to keep the wires AWAY from
the hot plate!
7. When the KNO3 is completely dissolved, use a test-tube holder to move the test tube and the
temperature probe from the water bath to the computer, and plug the temperature probe in to
channel 1. Click
. Hold the test tube up to the light to look for the first sign of
crystallization. At the same time, stir the solution with a slight up and down motion of the
temperature probe. At the moment crystallization starts to occur, click
. Type the
solubility value in the edit box from column 3 in the table above (g per 100 g H2O) and press
ENTER. Note: if you click ‘Keep’ to soon, you can press cancel. After you have saved the data
pair, return the test tube to the test tube rack and place the temperature probe in the water bath for
the next trial. Do NOT press Stop .
8. Repeat Steps 6 and 7 for each of the other three test tubes, hitting
point. Here are some suggestions to save time:
to keep each data
•
One lab partner can be stirring the next KNO3-water mixture until it dissolves while the other
partner watches for crystallization and enters data pairs using the computer.
•
Test Tubes 1 and 2 may be cooled to lower temperatures using cool tap water in a 250-mL
beaker. This drops the temperature much faster than air. If the crystals form too quickly, briefly
warm the test tube in the hot-water bath and redissolve the solid. Then repeat the cooling and
collect the data pair.
Stop .
9. When you have finished collecting data for all 4 test tubes, click
Record the
temperature (in °C) from the four trials in Data Table 1 on pg. 10. Discard the four solutions
in the waste container provided. Clean your probe with water and dry it gently.
10. Print the graph only: File → Print Graph. Enter your name(s) and the number of copies you
want to print.
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Part 2: Effect of the van’t Hoff Factor on the Freezing Point of a
Solution
When a solute is dissolved in a solvent, the freezing temperature is lowered in proportion to the
number of moles of solute added. This property, known as freezing-point depression, is a colligative
property; that is, it depends on the ratio of solute and solvent particles, not on the nature of the
substance itself. The equation that shows this relationship is:
ΔTideal = Kf•m•i
Equation 1
where ΔT is the freezing point depression, Kf is the freezing point depression constant for a particular
solvent [1.86 (°C•kg)/mol for water in this experiment], m is the molality of the solution in (mol
solute)/(kg solvent), and i is the van’t Hoff factor. The value of i represents how many particles a
substance creates in solution. If a substance does not break apart (dissociate) in solution then the
numerical value of i = 1. This is true for soluble covalent/molecular compounds. If the substance
dissociates, then the ideal value of i equals the number of particles created in solution.
In this experiment, you will need to recall the freezing point of pure water. You will then add a
known mass of sodium chloride, NaCl and sodium phosphate, Na3PO4, solutes to known volumes of
distilled water, and determine the lowering of the freezing temperature of each of the two solutions
separately. By knowing the freezing point depression constant for water, calculating the molality of
each solution, as well as the ideal van’t Hoff factor for each solute, you will calculate the freezing
point depression, ΔTcalc, and compare it to the ideal ΔT using Equation 1.
MATERIALS
Computer
Vernier computer interface
Logger Pro
2 Temperature Probes
400-mL beaker
2 18 X 150-mm test tube
Rock salt
Ice
Sodium chloride
Sodium phosphate dedecahydrate
Thermometer
Read all instructions before starting. Be careful with water and solutions next to electrical
connections.
GENERAL PROCESS:
This experiment you will use two temperature probes and measure the difference in the freezing point
between pure water and a solution containing water as a solvent. You will measure the freezing point
of two solutions, one solution containing NaCl as the solute, and one solution containing Na3PO4•12
H2O as the solute. Each test tube will be cooled so that solvent crystallization (freezing) will occur.
Graphs will be generated that will show temperature vs. time that you will use to find the difference
in the freezing temperatures of each of two solutions. You will then compare these freezing points to
the freezing point of pure water (0.00°C).
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PROCEDURE
1. Obtain and wear goggles.
2. Prepare the computer for data collection by opening the Vernier software program and
LoggerPro. Then, use the following pathway: File  Open  Experiments  Chemistry
With Computers  09 Evaporation. You will need to confirm the temperature probes you are
using in the dialog box. You will need to adjust the set-up by selecting Experiment → Data
Collection. Choose the Collection tab, and change the Length to 12 minutes. In the same dialog
box, change the Sampling Rate to 6 samples/minute in the left-hand box. The vertical axis of the
graph should have temperature scaled from -5.0°C to 30.0°C. The horizontal axis should have
time scaled from 0 to 12 minutes.
3. At your bench-top, obtain a 400-mL beaker. Cover the bottom
of the beaker with a layer of rock salt, then add a layer of ice,
and then, finally, another sprinkling of rock salt. Fill up the
beaker in this manner until 2/3 of the height of the beaker is
reached, then fill up the beaker half-way with tap water. Note:
It may be necessary to add ice to the bath and decant any water
that comes too close to the top of the beaker. At all costs avoid
spilling any water on the electronic equipment! Have an empty
beaker available to hold the excess water. You will use this
beaker of ice at a computer station.
4. Obtain two test large test tubes and fill each up with 10.0 grams of distilled water. (note: you will
need to pre-weigh a beaker and the test tube, tare it on the scale, and then add water until you
have weighed out 10 grams of water.) Record the exact amount of water in Data Tables 2 and 3
on pgs. 12 and 13. Weigh out enough NaCl and Na3PO4•12 H2O to make 0.20 molal solutions of
each, recording the exact amounts in Data Tables 2 and 3. (Note: you determined these amounts
on your pre-lab) Place each salt in its correct test tube (you may label them with tape to ensure
you know which salt is in each test tube) and thoroughly mix using a clean and dry stirring rod. If
you have difficulties getting the solids to dissolve, ask the lab instructor for assistance.
5. Clean your two temperature probes with distilled water and carefully dry each with a kimwipe.
Place the first temperature probe in the test tube containing NaCl and the second temperature
probe in the test tube containing Na3PO4•12 H2O and place both test tubes in the rock salt/ice
water bath and press the collect key. Do not let the probes rest on the bottom of the test tubes.
While the probes are in the solutions, mix the solutions swiftly with the probes in a circular
motion. You should scrape the sides of the test tubes with the probes while stirring. Once the
solution solidifies, stop stirring as you could damage the probes.
Note: If the temperature probes do not move freely, do not attempt to pull the probes out of the
solid. When you have completed the experiment, unplug the probe at the interface and carry it
to the hot water bath to heat it to liquid again. Be careful so you do not heat the wire.
6. Continue with the experiment until curves level out. Continue obtaining data for 2-3 more
minutes. Once this has happened, you may press stop.
7. Use the hot water bath to melt the probe out of the two samples. Do not attempt to pull the
probes out—this will damage them. Carefully clean each probe with distilled water and wipe any
dry with a kimwipe.
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8. Clean up your rock salt bath. Carefully decant the water away from the remaining rock salt, and
place the remaining rock salt in the Rock Salt Collection Container.
9. Determine the freezing points of the two solutions. To determine the freezing temperature of the
solutions, you need to determine the mean (or average) temperature in the portion of graph with
nearly constant temperature. Move the mouse pointer to the beginning of the graph’s level region
for each substance. Press the mouse button and hold it down as you drag across the level region of
the curve, selecting only the points in the plateau. Click on the Statistics button, . The mean
temperature value for the selected data is listed in the statistics box on the graph. Record these
values as the freezing temperature in the Data Tables 2 and 3 on pgs. 12 and 13.
10. Re-title the graph appropriately, and print your graph window only. By hand, label each curve
with the freezing point and identity of the substance.
Part 3: Using Freezing-Point Depression to Find Molecular Weight
In this experiment, you will first find the freezing temperature of the pure solvent, lauric acid,
CH3(CH2)10COOH. You will then add a known mass of benzoic acid solute, C6H5COOH, to a known
mass of lauric acid, and determine the lowering of the freezing temperature of the solution. By
measuring the freezing point depression, ΔT, and the mass of benzoic acid, you can use the formula
above to find the number of moles and therefore molecular weight of the benzoic acid solute, in
g/mol.
Figure 2
MATERIALS
Computer
Vernier computer interface
Logger Pro
Temperature Probe
400-mL beaker
ring stand
Exp #4 Solutions
utility clamp
18 X 150-mm test tube
lauric acid
benzoic acid
Thermometer
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Read all instructions before starting. Be careful with water and solutions next to electrical
connections.
General Process:
This experiment you will use two temperature probes and measure the difference in the freezing point
of a pure substance, lauric acid, and the substance with benzoic acid as a solute. The process will be
almost identical to the one used for the cooling curve for sodium thiosulfate in experiment 2 this
quarter. Each test tube will be heated so that the material is a liquid, then you will monitor the cooling
process. The graph will show temperature vs. time and you will find the difference in the freezing
temperatures of the pure solvent, lauric acid, and the mixture.
PROCEDURE
1. Obtain and wear goggles.
2.
Prepare the computer for data collection by opening the Vernier software program and
LoggerPro. Then, use the following pathway: File  Open  Experiments  Chemistry
With Computers  09 Evaporation. You will need to confirm the temperature probes you are
using in the dialog box. You will need to adjust the set-up by selecting Experiment → Data
Collection. Choose the Collection tab, and change the Length to 600 seconds. The vertical axis
of the graph should have temperature scaled from 0°C to 100°C. The horizontal axis should have
time scaled from 0 to 600 seconds.
3. Next to the computer, place a ring stand with either two utility clamps or a buret clamp on it to
hold the test tubes.
4. Obtain a test tube containing pure lauric acid and one containing lauric acid plus benzoic acid.
These test tubes are already prepared. In Data Table 4 on pg. 14, record the precise amounts of
lauric and benzoic acid in the mixture, from the label on the test tube.
5. Using a warm water bath on the lab counter, melt the two substances in the two test tubes. You
will need to stir a lot. You may use the temperature probes to gently stir the solutions. Unplug
the probes at the connection to the interface, but do not unplug the interface from the computer.
Do not contaminate one sample with the other.
6. When both are melted move them, in the warm water bath, to the ring stand by your computer.
Plug the temperature probes back into the interfaces, noting which is probe 1 and which is probe
2. About 30 seconds are required for the probe to warm up to the temperature of its surroundings
and give correct temperature readings. Both samples should be over 50°C. Quickly remove the
test tubes from the warm water bath and clamp them to the ring stand. Then click Collect to begin
data collection.
7. Using a swirling motion with the Temperature Probes, continuously stir the lauric acid or lauricbenzoic mixture during the cooling. Hold the top of the probe and not its wire. If the solutions
begin to get “sticky”, stop stirring or you will damage the probes.
Note: do not attempt to pull the probe out of the solid. When you have completed the
experiment, unplug the probe at the interface and carry it to the hot water bath to heat it to
liquid again. Be careful so you do not heat the wire (see step 9).
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8. Continue with the experiment allowing the data to collect for an additional 10 minutes. Click
Stop to end data collection.
9. Unplug the temperature probes from the interface and use the hot water bath to melt the probes
out of the two samples. Be sure to secure the wires so that they do not touch the water bath or the
hot plate! Do not attempt to pull the probe out—this will damage it. Carefully wipe any excess
sample liquid from the probe with a paper towel or tissue, and then wash the temperature probes
with soap and water. Dry them gently. Return the test tube containing lauric acid and lauricbenzoic mixture to the reagent bench.
10. To determine the freezing temperature of pure lauric acid, you need to determine the mean (or
average) temperature in the portion of graph with nearly constant temperature. Move the mouse
pointer to the beginning of the graph’s flat part. Press the mouse button and hold it down as you
drag across the flat part of the curve, selecting only the points in the plateau. Click on the
Statistics button, . Select the temperature which corresponds to the pure Lauric Acid, click
“OK.” The mean temperature value for the selected data is listed in the statistics box on the
graph. Record this value as the freezing temperature of lauric acid in Data Table 4. Click on the
upper-left corner of the statistics box to remove it from the graph.
11. When you have completed Step 10, click on the Examine button, . To determine the freezing
point of the benzoic acid-lauric acid solution, you need to determine the temperature at which the
mixture initially started to freeze. Unlike pure lauric acid, cooling a mixture of benzoic acid and
lauric acid results in a gradual linear decrease in temperature during the time period when
freezing takes place. (See Figure 3) As you move the mouse cursor across the graph, the
temperature (y) and time (x) data points are displayed in the examine box on the graph. Locate
the initial freezing temperature of the solution, as shown here - it may not be a distinct point, so
you will need to extrapolate the point. Record the freezing point in Data Table 4.
12. Re-title the graph appropriately, and print your graph window. By hand, label each curve with the
freezing point and identity of the substance.
f.p. of pure
f.p. of mixture
Figure 3
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SOLUTIONS
Fill-in
Name ____________________________
Prelab attached (p15)
Partner ___________________________
Stamp Here
Lecturer
Date
Part 1: Effect of Temperature on the Solubility of a Salt
DATA TABLE 1: SOLUBILITY OF KNO3 (must be completed in lab)
Trial
Solubility
(g / 100 g H2O)
1
40.0
2
80.0
3
120.0
4
140.0
Temp
(°C)
Make sure to fill out the temperatures before closing LoggerPro
Processing the Data
(This section may be completed in the lab or at home. It is not required for the final stamp.)
1. Draw a best-fit curve, by hand, for your data points on the printed graph. Label which side of the
graph corresponds to a saturated solution and which side corresponds to an unsaturated solution.
2. According to your data, how is solubility of KNO3 affected by an increase in temperature of the
solvent? (circle the answer that best applies)
Increases
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Decreases
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Stays the same
Page 10 of 15
3. Label the following points (as 3a, 3b and 3c) on your solubility graph. Using your printed graph,
determine if each of these solutions would be saturated or unsaturated:
a. 110 g of KNO3 in 100 g of water at 40°C
___________________________
b. 60 g of KNO3 in 100 g of water at 70°C
___________________________
c. 140 g of KNO3 in 200 g of water at 60°C
___________________________
4. According to your graph, will 50 g of KNO3 completely dissolve in 100 g of water at 50°C?
Label this point as the #4 on your graph.
5. According to your graph, will 120 g of KNO3 completely dissolve in 100 g of water at 40°C?
Label this point as the #5 on your graph.
6. According to your graph, about how many grams of KNO3 will dissolve in 100 g of water at
30°C? Label this point as the #6 on your graph.
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Part 2: Effect of the van’t Hoff Factor on the Freezing Point of a
Solution
Put calculated answers in SCIENTIFIC NOTATION! Watch Sig Figs!
DATA TABLE 2 FOR NaCl SOLUTION: (To obtain your final stamp, you need to acquire items 1, 3 and 6)
1) Grams of water in the test tube
g
2) Kilograms of water (solvent) in the test tube
kg
3) Mass of NaCl actually added to the test tube (remember this amount
should be close to the amount calculated on your pre-lab)
g
4) Molality of NaCl solution (SHOW WORK)
mol NaCl/kg H2O
5) Textbook (whole number) van’t Hoff factor value for NaCl
iideal =
6) Freezing point of the NaCl solution
o
C
7) ∆Tcalc calculation: (Tfreezing point of pure water – TNaCl solution)
o
C
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DATA TABLE 3 FOR Na3PO4•12 H2O SOLUTION: (To obtain your final stamp, you need to acquire items
1, 3 and 6)
1) Grams of water in the test tube
g
2) Kilograms of water (solvent) in the test tube
kg
3) Mass of Na3PO4•12 H2O actually added to the test tube (remember this
amount should be close to the amount calculated on your pre-lab)
g
4) Molality of Na3PO4•12 H2O solution (SHOW WORK)
mol Na3PO4/kg
H2 O
5) Textbook (whole number) van’t Hoff factor value for NaCl
iideal =
6) Freezing point of the NaCl solution
o
C
7) ∆Tcalc calculation: (Tfreezing point of pure water – TNaCl solution)
o
C
Questions:
1. Which solution had the larger ∆T value? ___________________________
2. Given that the concentrations of both solutions were the same, what variable determines the
magnitude of ∆T?
__________________
3. The theoretical ∆T for NaCl is 0.74°C. How does this compare with your calculated ∆T for the
NaCl solution? Calculate a percent difference for your ∆T value, and report this percent
difference to three significant figures. Show your calculation.
4. The i value assumes that an ionic compound dissociates 100% in solution. Is this true? Explain
this by considering the percent difference in ∆T values and the behavior of ions in solution (you
might want to read up on ion behavior in Tro, page 554).
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Part 3: Using Freezing-Point Depression to Find Molecular Weight
DATA TABLE 4 FREEZING POINT DEPRESSION (must be completed in lab)
Put calculated answers in SCIENTIFIC NOTATION! Watch Sig Figs!
Mass of lauric acid in mixture
g
Mass of benzoic acid in mixture
g
Freezing temperature of pure lauric acid
°C
Freezing point of the benzoic acid–lauric acid mixture
°C
Freezing temperature depression, ∆T: comparing pure and mixture
°C
Molality of the solution (ΔT = Kf • m• i and Kf = 3.9°C•kg/mol)
Show calculations
mol/kg
The solvent is:
The solute is:
Calculate the moles of benzoic acid from the molality
Show calculation
mols
Calculate the molar mass of benzoic acid (experimental)
Show calculation
g/mol
Structure
of
benzoic
acid
Molar mass
of benzoic
acid
(actual)
g/mol
Percent difference Show calculation
%
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Stamp here
SOLUTIONS
Prelab Questions
1. The solubility of sodium chloride, NaCl is 36g/100g H2O at 20°C. Indicate whether the resulting
solutions of sodium chloride in water are saturated, unsaturated, or supersaturated.
i. A stockroom technician dissolves 5.8 g of NaCl in 100mL of water at 20°C for a conductvity
test solution.
__________________
ii. 48.0 g of NaCl are mixed with 116.7 g of water, but does not dissolve completely.
__________________
iii. At 40°C 20.0 g of NaCl is dissolved in 50.0 g water then cooled back down to 20°C with the
salt remaining in solution.
_________________
2. A solution containing 2.00 g of an unknown substance in 50.0 g of cyclohexane freezes at 1.05°C.
The normal freezing point. of cyclohexane is 6.60°C and Kf = 20.4 °C/m. Both substances are
molecular.
a) Calculate the molality of the solution using ΔT = Kf•m•i
b) The units of molality are moles/kg. Moles of what over kilograms of what? Be more specific
than solute and solvent – label the species!
Moles of
Kilograms of
c) How many moles of the unknown substance are in this solution?
d) What is the molar mass of the unknown substance?
3. Calculate the number of grams of NaCl and Na3PO4•12 H2O you will need to create a 0.20 molal
solution using 10.00 grams of solvent (water). (use the back of this page for the calculations)
NaCl = 58.44 grams/mol
Mass of NaCl
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Na3PO4•12 H2O = 380.18 grams/mol
Mass of Na3PO4•12 H2O
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