This experiment studied the thermal decomposition of potassium chlorate (KClO3) in the absence of a catalyst. Three possible reactions were predicted to occur, producing either KCl, KClO2, or KClO. Oxygen gas produced was collected over water and its volume measured. The mass of the solid product was also measured and used to determine which reaction occurred. Analysis of the gas and solid products indicated that the decomposition reaction produced KClO2 in the absence of a catalyst, rather than KCl as occurs with a catalyst present.
1.Dew Point with non-condensable components
2.Flash with liquid vapor products
3.Condenser and Flash drum for ammonia synthesis
4.Azeotrope
Ideal Solutions vs. Azeotropes
Types of Azeotropes
• Number of Constituents:
• Heterogeneous or Homogeneous:
• Positive or Negative:
5.Enthalpy change of mixing
6.Solutropes
1.Dew Point with non-condensable components
2.Flash with liquid vapor products
3.Condenser and Flash drum for ammonia synthesis
4.Azeotrope
Ideal Solutions vs. Azeotropes
Types of Azeotropes
• Number of Constituents:
• Heterogeneous or Homogeneous:
• Positive or Negative:
5.Enthalpy change of mixing
6.Solutropes
A series of laws in physics that predict the behavior of an ideal gas by describing the relations between the temperature, volume, and pressure. The laws include Boyle's law, Charles' law, and the pressure law, and are combined in the ideal gas law
BC Chemistry 162 Laboratory Manual Experiment 6 Vapor Press.docxrosemaryralphs52525
BC Chemistry 162 Laboratory Manual
Experiment 6: Vapor Pressure of Liquids
- 1 -
Experiment 6: Vapor Pressure of Liquids
Background
Liquids contain molecules that have different kinetic energies (due to different velocities). Some of the
faster liquid molecules have enough kinetic energy to vaporize. At the same time, some of the slower
vapor molecules condense into liquid. In an open container, the rate of vaporization will be greater than
the rate of condensation—hence, the liquid will eventually evaporate. In a sealed flask, however, there
will be a point in which equilibrium is reached between the rate of vaporization and the rate of
condensation. To the eye, it seems that the liquid doesn’t change at equilibrium. But at the microscopic
level a vapor molecule enters the liquid phase for every liquid molecule that enters the gas phase.
The total pressure in the sealed flask is due to the vaporized liquid plus air molecules present in the flask:
Ptotal = Pvapor + Pair (1)
In this experiment, you will investigate the relationship between
the vapor pressure of a liquid and its temperature. Pressure and
temperature data will be collected using a gas pressure sensor and
a temperature probe (Figure 1). Vapor pressures will be
determined by subtracting atmospheric pressure from the total
pressure.
The flask will be placed in water baths of different temperatures to
determine the effect of temperature on vapor pressure. You will
measure the vapor pressure of methanol and ethanol and
determine the enthalpy (heat) of vaporization for each liquid.
Objectives
In this experiment, you will
Investigate the relationship between the vapor pressure of a liquid and its temperature.
Compare the vapor pressure of two different liquids at the same temperature.
Use pressure‐temperature data and the Clausius‐Clapeyron equation to determine the heat of
vaporization for each liquid.
Caution!
The alcohols used in this experiment are flammable and poisonous. Avoid inhaling their vapors. Avoid
contacting them with your skin or clothing. Be sure there are no open flames in the lab during this
experiment. Notify your teacher immediately if an accident occurs.
Procedure
1. Wear goggles! You will work in pairs for this lab, but you may share water baths with your table.
2. Prepare four water baths: 20 to 25°C (use room temperature water), 30 to 35°C, 40 to 45°C, and 50 to
55°C. You should also have some hot water on a hot plate on reserve.
3. Obtain a temperature probe and gas pressure sensor. The sensor comes with a
rubber‐stopper assembly (Figure 2). The stopper has three holes, one of which
is closed. Make sure your tubing and valve are not inserted in the closed hole.
Insert the rubber‐stopper assembly into a 125 mL Erlenmeyer flask.
Important: Twist the stopper into the neck of the flask to ensure a tight
fit.
Figure 1
Figure 2
BC Ch.
Armfield Gas Absorption Column ExperimentHadeer Khalid
The absorption of CO2 from air to water was studied in Gas absorption column built by Armfield company. Lab report and experiment was part of Separation Lab.
Chem 162 Lab 3: Gas Laws Part I & II- Sample Data for the class
1) Sample Data Group 1:
Part I
Part II
Volume (ml)
Pressure (kPa)
Temperature (°C)
Pressure (kPa)
103.0
60
70.8
113.5
88.0
70
66.3
112.6
73.0
85
61.8
111.5
62.0
100
57.1
110.4
44.0
140
51.5
109.0
34.0
180
39.9
105.5
31.0
200
26.4
101.8
10.5
96.7
2) Sample Data Group 2:
Part I
Part II
Volume (ml)
Pressure (Torr)
Temperature (°C)
Pressure (kPa)
32.0
630
57
109.6
29.2
690
52
108.4
27.8
726
48.5
107.4
25.6
790
43.6
106.3
24.2
843
38.1
104.8
22.2
914
33.1
103.5
29.3
102.2
25.4
101.1
22.5
100.1
20
99.4
17.4
98.6
12.8
97.2
9.4
96.7
Bellevue College | Chemistry 162
1
Empirical Gas Laws (Part 3): The Ideal Gas Law
Determination of the Universal Gas Constant, R
In this experiment, you will generate and collect a sample of hydrogen gas over water by the
reaction of magnesium with hydrochloric acid.
Using the Ideal Gas Law (PV=nRT) you will find values for the pressure (P), volume (V),
number of moles of the gas (n), and the temperature (T) in order to determine the gas constant
(R). Because there will be water vapor present in your sample, you will make a correction to the
measured pressure and then compare your result for R to the literature value.
In this experiment, you will:
Determine a value for the Universal Gas Constant, R. (Part 3 of Empirical Gas Laws)
Safety Precautions
Wear your goggles at all times. Hydrochloric acid is corrosive.
Avoid spills and contact with your skin and clothing. If HCl
comes in contact with your skin, inform your teacher and flush
the acid with large quantities of water.
Note: If you are doing Part 3 to determine the value of the Universal
Gas Constant, R in the same period as Parts 1 and 2, you should get Part 3
started first.
EXPERIMENTAL PROCEDURE (WORK IN PAIRS)
1. Put on goggles. Keep them on during the entire experiment.
2. Obtain a piece of magnesium ribbon that weighs no more than 0.08 grams. Record the mass
obtained (use significant figures!). Record this value in your data table (see report sheets).
Loosely roll it into a ball or coil it.
Encase the magnesium in a piece of copper mesh. Why do you think this might be helpful?
3. Fill the 800-mL beaker with approximately 200-mL of tap water.
4. Fill the 100-mL graduated cylinder with tap water. Using parafilm, a one-
hole stopper, or the palm of your hand, cover the top and invert the cylinder
into the beaker of water. You will end up with an inverted cylinder full of
water. Remove the parafilm or stopper if you used one. Rest the cylinder
on the bottom of the beaker. Try not to introduce any air bubbles in your
inverted cylinder (see Figure 1).
5. Place the magnesium (in its copper cage) into the graduated cylinder. Make
sure the magnesium is captured in the cylinder.
Figure 1: Gas collection in an
inverted cylinder full of water.
Lab 9 Chemical Reactions IIPre-lab Questions1. What is a limi.docxsmile790243
Lab 9: Chemical Reactions II
Pre-lab Questions
1. What is a limiting reagent?
2. A student used 7.15 g of CaCl2 and 9.25 g of K2CO3 to make CaCO3. The actual yield was 6.15 g of CaCO3. Calculate the limiting reagent and the percent yield.
Experiment: Synthesis of Garden Lime
Procedure
**Take photographs of your experiment set up and your results. Submit them with your laboratory report.**
1. Table 1 provides an example set of data for 1.0 g CaCl2.
2. For Trial 1, weigh into a 250 mL beaker the amount of calcium chloride (CaCl2) shown in Table 1. Record the exact mass you weigh out in the Trial 1 column of the Data section.
3. Measure 50.0 mL of distilled water into a 100 mL graduated cylinder. Pour the water into the 250 mL beaker with the calcium chloride.
4. Stir the solution with a stirring rod until all of the calcium chloride is dissolved.
5. Weigh out 2.5 g of potassium carbonate (K2CO3) in a 50 mL beaker. Record the exact mass in the Data section.
6. Measure 25.0 mL of distilled water into a 100 mL graduated cylinder. Add the water into the 50 mL beaker containing the potassium carbonate.
7. Stir the potassium carbonate in the distilled water with a stirring rod until it is all dissolved.
8. Pour the K2CO3 solution into the 250 mL beaker that has the CaCl2 solution. Rinse the beaker that contained the K2CO3 with a few mL of water and add this to the CaCl2 solution. Stir the mixture.
9. As soon as the reaction begins, record your observations in the Data section. Continue stirring until you see no more precipitate forming.
10. Set up the funnel in the Erlenmeyer flask as shown in Figure 2.
HINT: Do NOT begin filtering yet!
11. Zero the scale and weigh a piece of filter paper and a watch glass. Record the masses of both items in the Data section.
12. Prepare a filtering funnel as shown in Figure 2: fold a piece of filter paper in half twice to make quarters, and open the paper to make a small cone (three quarters are open on one side and one quarter is on the opposite side). Place the paper cone into the funnel and hold it in place with your fingers. Pour a small amount of distilled water through the paper to secure it inside the funnel.
13. Filter the mixture by pouring it into the filter paper in the funnel. Use the stirring rod and distilled water in a wash bottle to transfer the entire solid into the filter paper.
HINT: For best results, be sure to transfer all of the precipitate into the filter paper. Use a rubber policeman if it is available to help with the transfer.
14. Rinse the remaining solid in the filter paper twice with distilled water from a wash bottle to rinse off excess sodium chloride (NaCl). After all the liquid has filtered through, rinse the product with approximately 5 mL of ethanol to aid in its drying. Allow the ethanol to completely finish filtering through the paper.
15. Remove the filter paper carefully so as to not lose any product. Gently unfold the filter paper and lay it flat on the pre-weighed wat ...
1. The Decomposition of Potassium Chlorate
Small quantities of molecular oxygen (O2) can be obtained from the thermal decomposition of
certain oxides, peroxides, and salts of oxoacids. Some examples of these reactions are
2 Ag2O(s) 4 Ag(s) + O2(g)
2 BaO2(s) 2 BaO(s) + O2(g)
MnO2
2 KClO3(s) 2 KCl(s) + 3 O2(g)
The last reaction, the decomposition of potassium chlorate, includes manganese(IV) oxide
(MnO2) as a catalyst. A catalyst is a substance that causes an increase in the rate of a chemical
reaction without being used up in the reaction. It is this reaction that will be studied in this
experiment.
The thermal decomposition of KClO3 (potassium chlorate) in the absence of a catalyst will be
studied. The identify of the solid that remains after the decomposition can be determined from
the quantity of oxygen that is evolved. Identification can then be made by comparing the
measured mass of the solid product with a calculated value based on the quantity of O2.
The solid product that results from the thermal decomposition of KClO3 is KCl when the catalyst
MnO2 is present. What will happen if the catalyst is not added? Will the loss of oxygen be less
extensive? If so, the solid product could be either KClO2 (potassium chlorite) or KClO
(potassium hypochlorite). This experiment will investigate these questions. The three reactions
that could possibly occur can be written and balanced based on these predictions.
A sample of KClO3 of known mass will be heated in the absence of a catalyst until the evolution
of oxygen is complete. Oxygen will be collected in a flask by the displacement of water. The
volume of water displaced equals the volume of O2 gas produced. In order to determine the
correct stoichiometry of this reaction, you will need to obtain the number of moles of O2 that
have been evolved. You can calculate this quantity from the rearranged form of the ideal gas
law:
RT
PV
n =
where P refers to the partial pressure of oxygen in the collected gas mixture, V is the volume of
water displaced, T is the Kelvin temperature of the gas mixture, and R is the constant. A
commonly used value for R is 0.082056 L·atm/mol·K. If this value for R is used, then P must be
expressed in atmospheres and V in liters.
Since the oxygen is collected over water, water vapor will also be present in the gas. The
experiment is designed so that the total pressure of the oxygen and water vapor will be equal to
the atmospheric pressure:
2. OHOTotal 22
PPP +=
you can easily measure atmospheric pressure with a barometer. The partial pressure of oxygen
in the flask can be calculated by subtracting the vapor pressure of water from the
atmospheric pressure. Table I gives the vapor pressure of water at various temperatures.
Figure I shows the apparatus for this experiment. The sample of KClO3 is placed in the test tube
and the Erlenmeyer flask is filled with water. Some of the water is displaced by oxygen and is
pushed into the beaker. The volume of water in the beaker will be identical to the volume of
oxygen in the flask.
Procedure
1. Record the atmospheric pressure from the laboratory barometer. This will equal the
pressure of the (O2 + H2O) gas mixture that develops in your Erlenmeyer flask.
2. Caution: KClO3 is a very strong oxidizing agent. Make certain you place the lid
back on the bottle containing the KClO3 after you obtain your sample. Do not let
this substance contact paper or the rubber stopper in the test tube of the apparatus.
3. Record the weight of the beaker on a top-loading balance. Assemble the apparatus as
shown in Figure I.
4. Fill the Erlenmeyer flask with distilled water, so that the level of the water is about 1 inch
below the short glass tube. Open the pinch clamp. Remove the stopper from the test
tube. Use a suction bulb to force air through the glass tube (test-tube end) until the
rubber tube is filled with water. Allow a little water to enter the beaker to about 2 inches.
Close the pinch clamp near the end of the tubing where the water will exit.
5. Make sure that the test tube is clean and dry. Take the test tube and a clean, empty, dry
400 mL beaker to the top-loading balances. Tare the empty beaker, add the test-tube and
record its mass. Tare the beaker and test-tube.
6. Carefully add a small amount of KClO3 to the test tube. Continue to add KClO3 until you
have about 1.0 g of KClO3 in the test tube. A sample in the range of 0.9 g to 1.1 g of
KClO3 will work. Record the mass to 0.01 g.
7. Clamp the test-tube to the ringstand and stopper the test tube.
8. Open the pinch clamp. Lift the beaker with your hands until the water level in the beaker
is equal to the water level in the Erlenmeyer flask. When the water levels are equal, have
your partner close the pinch clamp. This equalizing process will ensure that the pressure
acting on the water in the beaker (atmosphere) is equal to the pressure acting on the water
in the flask.
3. 9. Empty the beaker but do not dry it. The volume of the water drops that remain in the
beaker will be roughly equal to the volume that will remain after the displaced water is
poured into a graduated cylinder for measurement.
10. Place the glass tube (connected to the hose) back into the beaker. MAKE CERTAIN
THE PINCH CLAMP IS OPEN!
Caution: If the clamp is not opened at this point, the build-up of gas during heating could
cause an explosion, although it is more likely that a stopper would be forced to loosen.
Also, make certain that the longer glass rod is not touching the bottom of the Erlenmeyer
flask. This would also result in a closed system and an explosion could result.
11. Heat the test tube. The solid will melt, oxygen will be evolved, and water from the flask
will be displaced into the beaker. Be cautious at first and brush the flame over the test
tube. After a few minutes when the liquid solidifies, the test tube can be heated more
strongly. One gram of KClO3 reactant should cause the displacement of between 250 and
300 mL of water.
12. Heat the solid thoroughly until no more gas is evolved. The contents of the test tube will
solidify, since the melting point of the product is greater than that of KClO3.
13. Turn off the flame and allow the system to come back to room temperature. Allow five
minutes for this process.
14. As in step 8, equalize the water levels (this may require lifting the Erlenmeyer flask) and
close the clamp.
15. Remove the tube from the beaker. Record the mass of the beaker plus the water on a top-
loading balance. Use Table I to determine the volume of water displaced.
16. Measure the temperature of the water to the nearest degree. Assume this is the
temperature of the gas. Determine the appropriate vapor pressure of water from Table II.
17. Obtain the mass of the test-tube and its contents. Calculate and record the mass of the
product.
18. If you have any reason to believe that your experiment did not work (e.g. you did not get
ca. 300 mL of water transfer during the experiment), repeat steps 6-24 with a second
sample of KClO3. If you believe that you were able to capture the complete volume of
oxygen gas, a second trial is not necessary.
4. Figure I: A diagram of the laboratory setup used in this experiment.
Table I Density (g/mL) of Water at Various Temperatures (o
C)
Temp. Density Temp. Density Temp. Density
17 0.9988 22 0.9978 27 0.9965
18 0.9986 23 0.9976 28 0.9962
19 0.9984 24 0.9973 29 0.9959
20 0.9982 25 0.9971 30 0.9956
21 0.9980 26 0.9968 31 0.9953
Table II Vapor Pressure of Water (torr) (data from the CRC Handbook of Chemistry
and Physics, 49th
edition, 1968.)
Temp. Vapor Pressure Temp. Vapor Pressure Temp. Vapor Pressure
17 14.5 22 19.8 27 26.7
18 15.5 23 21.1 28 28.3
19 16.5 24 22.4 29 30.0
20 17.5 25 23.8 30 31.8
21 18.7 26 25.2 31 33.7
5. Questions
1. Write a balanced chemical equation for each of the three possible reactions that could
occur when potassium chlorate, KClO3 is thermally decomposed.
2. Suppose the atmospheric pressure when you performed your experiment was 751 torr.
The temperature of your water was found to be 23.0°C. What is the pressure of the
oxygen gas, 2OP , that is produced?
Data Treatment
1. Determine:
a. moles of oxygen that were produced.
b. moles KClO3 that were reacted.
c. ratio of the number of moles of oxygen to the number of moles of KClO3.
Write the chemical equation for the reaction that the data suggest occurred.
2. Using the mass of reactant KClO3 that you used in your most successful trial, calculate
how many grams of each of the solid products of each possible reaction would
theoretically be produced in the decomposition. Which decomposition reaction occurred,
based on the mass of solid generated in the reaction? Write the chemical equation for
the reaction that the data suggest occurred:
3. You determined which chemical reaction occurred by analyzing the amount of gas
produced in the reaction, and the mass of solid product remaining after reaction. Which
determination method do you believe is more reliable, the analysis of the mass of the
solid product or the analysis of the gas produced? Explain your answer.
4. In the presence of a catalyst, KCl is produced in the thermal decomposition of KClO3.
You did not use a catalyst. Did you get the same products for this reaction as you would
have with a catalyst?