Experiment Freezing Point of Stearic Acid _Formal Lab report_ by niusheng11

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Experiment Freezing Point of Stearic Acid (Partner 1
do this) (Formal Lab report-email)
Purpose: To make a temperature/time graph to show potential and kinetic energy changes as stearic acid
cools from a solid to a liquid. The stearic acid will be in a test tube in a beaker in the beaker cabinet and will
have a thermometer in it. When you finish put this apparatus back in the cabinet.
1. Put some stearic acid into a large test tube, and heat gently until it all melts. Let the stearic acid cool to
around 90°C before you put the thermometer into the stearic acid, or the thermometer will break when you
put it into the melted stearic acid.
2. Record the temperature every 15 seconds as the stearic acid cools, until it solidifies.
3. Remove the thermometer from the stearic acid, and wash it, but do not attempt to clean the stearic acid out
of the test tube. Merely parafilm (Back counter top by black hood) the stearic acid test tube so it can be used
in the future.
4. You can use the computer to make a graph with temperature on the y axis and time on the x axis or use
graph paper.
Determine freezing point of stearic acid based on your graph. Calculate your percent error. Look up the
freezing point of stearic acid in the CRC (Maroon book on book shelf at the back of the room).

Melting Point of Naphthalene (Partner 2 do this-email)
The Lab Purpose: Determine the melting point temperature of naphthalene?
Our Hypothesis: The "theoretical" melting point value should be your hypothesis. Use the reference
books at the back of the room to determine the melting point of naphthalene.
A hot water bath is suggested to heat the naphthalene slowly. Get a beaker 2/3 full of water and heat the water
until it is 95+ degrees. Remove the beaker from the hot plate and place the test tube with the naphthalene
into the hot water. Record the temperature change at 10 second intervals. When the temperature change
slows, take measurement every 5 sec. When you believe you have reached the melting point temperature,
continue heating slowly for another 60 seconds. Graph this data on a separate sheet of graph paper. Let the
test cool and replace in the rack on the center station.

FORMAL LAB REPORT: Data table, graph, % error, and conclusion. In the conclusion
express the significance of the graph shape (flat part)
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Experiment Boiling Water (Formal Lab Report-use Excel to graph data and
do a wiki page with the Excel embedded) both
Purpose: To see the boiling point plateau of water.
Place a beaker full of ice on the hot plate. Stir the ice/water in the beaker before taking the temperature. Use
a thermometer to take the temperature of water every 15 seconds until the water has boiled for at least 5 min.
Record the temperature every 15 seconds five minutes after the water has begun to boil. Caution: Be careful
not to burn yourself on the steam. You can use Excel to make a graph with temperature on the y axis and
time on the x axis or use graph paper. See me for an easy way to record the temperatures in Excel. Use the
graph to determine the melting point and boiling point of water. Compare these to the accepted value. Do a
percent error for the freezing and boiling point. Take the average temperature while the ice was melting and
the average temperature while the water was boiling. Change this average and the accepted value to kelvin
(K = oC + 273) before calculating the % error. Discuss the significance of the shape of the graphs (include
thoughts about why the graph show some sections with a slope of zero.)

    PHYSICAL PROPERTIES OF WATER:                                                   Both-Handwritten
Density And Superabsorbent Polymers
INTRODUCTION: Sodium polyacrylate is a nontoxic superabsorbent polymer found in high absorbency
diapers and marketed under the tradename Waterlock®. It is produced by polymerizing sodium acrylate and
acrylic acid. Osmotic pressure causes the superabsorbent polymer to absorb water in an effort to equilibrate
the sodium ion concentration inside and outside the polymer. The concentration of electrolyte in the water
being absorbed affects the amount of water the polymer can absorb.
PURPOSE: This experiment is intended as an introduction to measurement and graphing. You will use a
balance and a graduated cylinder to collect data. This data will then be plotted using Graphical Analysis.
PROCEDURE: Read through the lab to avoid having to redo the lab.
1. Weigh a clean, dry 10-mL graduated cylinder. Record this mass in the Data Table.
2. Weigh a dry 100-mL graduated cylinder and record the mass. Fill the 100-mL cylinder to the 100 mL
    mark with distilled water, weigh and record the mass.
3. Place a small amount of Sodium polyacrylate (about 0.5 mL in a 10-mL cylinder). Tap the base of the
    cylinder against the palm of your hand so the polymer is packed tightly with a flat upper surface.
4. Read the volume of the powder in the graduated cylinder. Remember, you need to estimate the last
    decimal place. So your volume will be recorded to the 100th of a milliliter.) Record this volume in the
    Data Table.
5. Determine the mass of the cylinder plus powder. Record.
6. Set the 10-mL graduated cylinder with the powder on a flat surface and add 0.5 - 1.0 mL of distilled
    water from the 100 ml graduated cylinder. (A disposable pipet may be used to avoid adding too much
    water at once.) The individual increments of water added from the 100-mL cylinder need not be
    measured. However, the water should be added quickly so that the top surface remains flat. Wait until the
    polymer gels. Read the volume of the polymer in the cylinder and record. Determine the mass of the
    cylinder with the polymer and record.
7. Repeat step 6 until the volume reaches approximately 10 mL. Be careful not to go over the 10 mL mark.
8. Obtain a 500 to 800 mL beaker. Using a spatula remove the superabsorbent polymer from the graduated
    cylinder and place in the 800-mL beaker. Wash any remaining polymer into the 800-mL beaker with
    small volumes of distilled water Note: This requires patience!
9. Determine how much water the polymer can hold by continuing to add increments of water until the
    polymer starts to flow when the beaker is tilted. Refill the 100 mL graduated cylinder with distilled water
    as necessary. Be sure to keep a record of the total amount of distilled water used. Use your experimental
    density to calculate the mass of the water added.
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10. If there are graduations on the side of the beaker estimate the total volume of the polymer. If the beaker is
    not graduated and a 1000-mL graduated cylinder is available, pour the polymer into a 1000-mL
    graduated cylinder and record the volume.
11. Use a Centigrade thermometer to measure the temperature of the distilled water. Record the temperature.
    Find the density of water at this temperature in the CRC Handbook Chemistry Physics) and record this
12. Add a pinch of salt to the beaker containing the polymer. Stir. Observe what happens. Record your
13. Pour the polymer and salt mixture into the large jar or beaker designated by your teacher for polymer
    disposal. Do not pour the polymer down the drain! Clean and dry your glassware and spatula.
14. Calculate the mass and the density of your polymer in steps 4--7. Determine how many times its volume
    in water your polymer can hold? Determine how many times its mass in water your polymer can hold?
15. Using Graphical Analysis III, plot the volume of the polymer as the x-axis and the mass of the polymer
    as the y-axis. Determine the slope of the line. You may use Excel to do your graphs.

QUESTIONS: Each partner does his/her own answers. DO NOT Work together.
1. How does the density of the polymer change as the water is added? Why?

2. What is the significance of the slope of the line? Explain. Calculate your percentage error. What point on
   your graph would you expect to be in greatest error? Why?

3. How many times its volume in distilled water can your polymer hold? How many times its mass in water
   can your polymer hold? Which of these two values is more accurate? Why? How does your calculated
   mass ratio compare with the claim that the polymer can hold about 800 times its weight in distilled

4. Are the water and the dry-polymer volumes additive in the production of the hydrated polymer?

5. Does the addition of salt cause a noticeable change in the consistency or in the volume of the polymer?

Data Table:
   A. Mass of empty 10 mL graduated cylinder                                                   g

   B. Mass of empty 100 mL graduated cylinder                                                  g

   C. Mass of cylinder and 100 mL of water (use .1 balance)                                    g

   D. Mass of 100 mL of water (C-B)                                                            g

   E. Density of water (experimental) [D/100 mL]                                               g
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     Volume of             Mass of Polymer   Mass of Polymer    Density of
     Polymer               & Cylinder                           Polymer










F. Slope of the line                    Show work:                            ________
G. Total volume of distilled water added                                      ________ mL
H. Mass of water added (G x E)                                                ________ g
I. Estimated volume of hydrated polymer in beaker                             ________ mL
or Volume of hydrated polymer measured in 1000-mL cylinder                    ________ mL
J. Temperature of distilled water                                             ________ oC
K. Density of water at the above temp. According to the CRC Handbook          ________g/mL
L. Density of water calculated in E                                           ________ g/mL
M. % Error [((L-K)/K)100] %                                                   ________ %
N. Slope of the line [F]                                                      ________ g/mL
0. % Error [((N-K)/K)100]                                                     ________ %
Q. Total amount of distilled water added [H, G]            ________ g         ________ mL
R. Amount of original polymer powder [#1]                  ________ g         ________ mL
Ratio of polymer to water by                         MASS _______ by VOLUME _______
ThermoLabs                                    Name: _________________________________ Page 5 of 11

Heat of Fusion of Ice Both                                 (Both-Handwritten)

      • Measure temperature changes when ice melts in water.
      • Calculate heat changes from experimental data.
      • Calculate the heat of fusion of ice from experimental data.
Consider this. You pop the top on the can of your favorite soft drink and pour it over ice. Does the soft drink
become cold simply because the ice is cold or is there more to it than that? You know that heat always flows
from a warm object to a cooler one. For example, when you place your hand on a cold window your hand
gets cold. Heat flows from your hand to the window giving your hand a cooling sensation. Similarly, when
you touch a hot stove heat flows from the stove to your hand and your hand feels hot. It makes sense then
that when a warm drink comes into contact with cold ice, heat will flow from the drink to the ice and the
drink will become cooler.
Something very important also happens: The ice melts! Melting ice is an endothermic process. That is, the
melting process absorbs heat from the surroundings causing the surroundings t cool off. It’s the melting ice,
not its cold temperature, that is most responsible for cooling the drink. The ability of ice to cool objects is
largely due to the fact that some of the ice melts when it comes into contact with warm objects.
The heat change associated with melting ice or any other solid can be expressed quantitatively. The amount
of heat absorbed from the surroundings when a specific amount of solid melts is called the heat of fusion.
Heats of fusion are usually expressed in the SI units of kilojoules per mole (kJ/mol). The heat of fusion of ice
is also commonly expressed in units of calories per gram (cal/g).
In this lab you will use the technique of calorimetry to measure the heat of fusion of ice, the amount of heat
absorbed when ice melts. You will place a piece of ice into some hot water and record the temperature
change of the system as the ice melts when energy flows from the hot water to the ice. You will first measure
the mass and temperature of the hot water. You will also need to know the mass and temperature of the ice
but because of possible significant heat losses you will determine these quantities indirectly. Because ice
melts at 0°C, it will suffice to assume that the temperature of the ice is 0°C. To minimize heat loss, the best
way to measure the mass of the ice is to measure the difference in mass of the hot water before and after you
place the ice into it. You will carry out the experiment in an insulated polystyrene coffee cup.
hot water (use the microwave)           alcohol thermometer                     polystyrene coffee cup
ice from the cafeteria                  one-ounce plastic cup                   balance

Experimental Procedure
1. Mass a polystyrene coffee cup.
2. Fill a one-ounce plastic cup 1/4 full with hot water from the tap. Transfer it to the tared coffee cup and
   weigh it. Record the weight in Table 14.1.
3. Measure the temperature of the water and record it in Table 14.1.
4. Place a small ice chip into the water, stir and record in Table 14.1 the lowest temperature reached after
   the ice melts completely.
5. Reweigh the entire system and record your results in Table 14.1.
6. Work through the calculations on the Experimental Data page and then repeat the experiment until your
   calculated heat of fusion of ice is consistent.
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Experimental Data: Record your results in Tables 14.1 and 14.2.
Table 14.1
           Quantity                Trial 1        Trial 2       Trial 3
a. Hot water weight (g)

b. Hot water temperature (oC)

c. Final Temperature (oC)

d. Total weight, hot water and ice (g)

e. Ice Temperature (oC)

Table 14.2
                     Quantity                       Calculation                Trial 1             Trial 2   Trial 3
f. Ice weight (g)                                        d-a

g. ΔT for hot water (oC)                                 c-b

h. Heat lost by water (cal) (1 is the number one)   a x (c – b) x 1

i. ΔT for ice (oC)                                       c -e

j. Heat gained by ice (cal) (1 is the number one)   f x (c – e) x1
l. Heat lost + heat gained (cal) Heat lost is
k. Heat of fusion of ice (cal/g)                         - l/f
                                                                 This is the letter L not a one.

Questions for Analysis
Use what you learned in this lab to answer the following questions.
1. Without first weighing the ice, how can you tell the weight of the ice from your data?

2. Why is ΔT for the hot water negative? What does the negative sign mean?

3. What does the negative sign of the heat lost by the water signify?
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4. Why can you assume the initial temperature of the ice is 0°C?

5. What is the sign of the heat absorbed by the ice? Explain.

6. The law of conservation of energy requires that the heat gained by the ice should balance the heat lost by
    the water assuming no losses of energy to the surroundings. Does item k in Table 14.2 of the calculated
    data equal zero? Explain.

7. Why does the heat required to melt the ice divided by the mass of the ice equal the heat of fusion of ice
    (heat of fusion = -j/f)? Why is if the negative sign necessary?

8. Given that one calorie equals 4.18 joules, use your experimental data to calculate the heat of fusion of ice
    in joules per gram.

9. Use your experimental data to calculate the molar heat of fusion of ice in kJ/mol.

11. Find the accepted value for the heat of fusion of ice in your textbook or the CRC (bookcase at the back of
    the room) and compare your calculated values with the accepted value. Calculate % error.
ThermoLabs                                    Name: _________________________________ Page 8 of 11

       Vapor Pressure and Heat of Vaporization
                                                 Email to me
When a liquid is placed in a container, and the container is sealed tightly, a portion of the liquid will
evaporate. The newly formed gas molecules exert pressure in the container, while some of the gas condenses
back into the liquid state. If the temperature inside the container is held constant, then at some point
equilibrium will be reached. At equilibrium, the rate of condensation is equal to the rate of evaporation. The
pressure at equilibrium is called vapor pressure, and will remain constant as long as the temperature in the
container does not change.
In mathematical terms, the relationship between the vapor pressure of a liquid and temperature is described
in the Clausius-Clayperon equation,
                                                      H vap  1 
                                            ln P              C
                                                        R T 
where ln P is the natural logarithm of the vapor pressure, ΔHvap is the heat of vaporization, R is the universal
gas constant (8.31 J/mol•K), T is the absolute, or Kelvin, temperature, and C is a constant not related to heat
capacity. Thus, the Clausius-Clayperon equation not only describes how vapor pressure is affected by
temperature, but it relates these factors to the heat of vaporization of a liquid. ΔHvap is the amount of energy
required to cause the evaporation of one mole of liquid at constant pressure.
In this experiment, you will introduce a specific volume of a volatile liquid into a closed vessel, and measure
the pressure in the vessel at several different temperatures. By analyzing your measurements, you will be
able to calculate the ΔHvap of the liquid.

In this experiment, you will
      Measure the pressure inside a sealed vessel containing a volatile liquid over a range of temperatures.
      Determine the relationship between pressure and temperature of the volatile liquid.
      Calculate the heat of vaporization of the liquid.

                                                     Figure 1

Vernier computer interface                  20 mL syringe
computer                                    two 125 mL Erlenmeyer flasks
Vernier Gas Pressure Sensor                 ethanol, CH3CH2OH
Temperature Probe                           400 mL beaker
rubber stopper assembly                     1 liter beaker
plastic tubing with two connectors          hot plate
 ThermoLabs                                    Name: _________________________________ Page 9 of 11
 1. Obtain and wear goggles. CAUTION: The alcohol used in this experiment is flammable and poisonous.
    Avoid inhaling the vapors. Avoid contact with your skin or clothing. Be sure that there are no open
    flames in the room during the experiment. Notify your teacher immediately if an accident occurs.

 2. Use a hot plate to heat ~200 mL of water in a 400 mL beaker.

 3. Prepare a room temperature water bath in a 1 liter beaker. The bath should be deep enough to completely
    cover the gas level in the 125 mL Erlenmeyer flask.

 4. Connect a Gas Pressure Sensor to Channel 1 of the Vernier computer interface. Connect a Temperature
    Probe to Channel 2 of the interface. Connect the interface to the computer with the proper cable.

 5. Start the Logger Pro program on your computer. Open the file “34 Vapor” from the Advanced Chemistry
    with Vernier folder.

 6. Use the clear tubing to connect the white rubber stopper to the Gas Pressure Sensor. (About one-half turn
    of the fittings will secure the tubing tightly.) Twist the white stopper snugly into the neck of the
    Erlenmeyer flask to avoid losing any of the gas that will be produced as the liquid evaporates (see Figure
    1). Important: Open the valve on the white stopper.

 7. Your first measurement will be of the pressure of the air in the flask and the room temperature. Place the
    Temperature Probe near the flask. When the pressure and temperature readings stabilize, record these
    values in the first column (Initial) of your data table.

 8. Condition the Erlenmeyer flask and the sensors to the water bath.
     a. Place the Temperature Probe in the room temperature water bath.
     b. Place the Erlenmeyer flask in the water bath. Hold the flask down into the water bath to the bottom of
        the white stopper.
     c. After 30 seconds, close the valve on the white stopper.
 9. Obtain a small amount of ethanol. Draw 3 mL of ethanol into the 20 mL syringe that is part of the Gas
    Pressure Sensor accessories. Thread the syringe onto the valve on the white stopper (see Figure 1).

10. Add ethanol to the flask.
     a.   Open the valve below the syringe containing the 3 mL of ethanol.
     b.   Push down on the plunger of the syringe to inject the ethanol.
     c.   Quickly pull the plunger back to the 3-mL mark. Close the valve below the syringe.
     d.   Carefully remove the syringe from the stopper so that the stopper is not moved.
11. Gently rotate the flask in the water bath for a few seconds, using a motion similar to slowly stirring a cup
    of coffee or tea, to accelerate the evaporation of the ethanol.

12. Monitor and collect temperature and pressure data.
     a.   Click         to begin data collection.
     b.   Hold the flask steady once again.
     c.   Monitor the pressure and temperature readings.
     d.   When the readings stabilize, click      .
13. Add a small amount of hot water, from the beaker on the hot plate, to warm the water bath by 3–5°C. Use
    a spoon or a dipper to transfer the hot water. Stir the water bath slowly with the Temperature Probe.
    Monitor the pressure and temperature readings. When the readings stabilize, click          .
 ThermoLabs                                      Name: _________________________________ Page 10 of 11
14. Repeat Step 13 until you have completed five total trials. Add enough hot water for each trial so that the
    temperature of the water bath increases by 3-5°C, but do not warm the water bath beyond 40°C because
    the pressure increase may pop the stopper out of the flask. If you must remove some of the water in the
    bath, do it carefully so as not to disturb the flask.

15. After you have recorded the fifth set of readings, open the valve to release the pressure in the flask.
    Remove the flask from the water bath and take the stopper off the flask. Dispose of the ethanol as

16. Click         to end the data collection. Record the pressure readings, as Ptotal, and the temperature
    readings in your data table.

17. Do not exit the Logger Pro program until you have completed 1–4 of the Data Analysis section.

                                  Initial     Trial 1     Trial 2      Trial 3      Trial 4    Trial 5

            Ptotal (kPa)

            Pair (kPa)

            Pvap (kPa)

            Temperature (°C)

 1. The Pair for Trials 2-5 must be calculated because the temperatures were increased. As you warmed the
    flask, the air in the flask exerted pressure that you must calculate. Use the gas law relationship shown
    below to complete the calculations. Remember that all gas law calculations require Kelvin temperature.
    Use the Pair from Trial 1 as P1 and the Kelvin temperature of Trial 1 as T1.
                                                       P1 P2
                                                       T1 T2
 2. Calculate and record the Pvap for each trial by subtracting Pair from Ptotal.

 3. Prepare and print a graph of Pvap (y-axis) vs. Celsius temperature (x-axis).
     a. Disconnect your Gas Pressure Sensor and Temperature Probe from the interface.
     b. Choose New from the File menu. An empty graph and table will be created in Logger Pro.
     c. Double-click on the x-axis heading in the table, enter a name and unit, then enter the five values for
        temperature (°C) from your data table above.
     d. Double-click on the y-axis heading in the table, enter a name and unit, then enter the five values for
        vapor pressure from your data table above.
     e. Does the plot follow the expected trend of the effect of temperature on vapor pressure? Explain.

 4. In order to determine the heat of vaporization, ΔHvap, you will first need to plot the natural log of Pvap vs.
    the reciprocal of absolute temperature.
     a. Choose New Calculated Column from the Data menu.
     b. Create a column ln vapor pressure.
     c. Create a second column, reciprocal of absolute temperature, 1/(Temperature (°C) + 273).
ThermoLabs                                    Name: _________________________________ Page 11 of 11
   d. On the displayed graph, click on the respective axes, and then select ln vapor pressure to plot on the y-
      axis, and reciprocal of absolute temperature to plot on the x-axis. Autoscale the graph, if necessary.
   e. Calculate the linear regression (best-fit line) equation for this graph. Calculate ΔHvap from the slope of
      the linear regression.
   f. Prepare and print a second graph.

5. The accepted value of the ΔHvap of ethanol is 42.32 kJ/mol. Compare your experimentally determined
   value of ΔHvap with the accepted value.

As you put together this lab, take a screen shot of each Graph and insert this into this file label each. Email
   to me.

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