Chemistry 125
Laboratory Manual
2004 - 2005
This packet contains the laboratory experiments for Chemistry 125. Please refer to your class syllabus
for the schedule of dates as to when the experiments will be performed. Please bring booklet with you
to each laboratory class session.
Table of Contents
Scientific Measurement .......................................................................................................3
Density ...............................................................................................................................13
Separation of a Mixture .....................................................................................................25
Determination of a Chemical Formula ..............................................................................31
Solubility Rules and Writing Formulas .............................................................................37
Acid Base Titration ............................................................................................................43
Cycle of Copper Reactions ................................................................................................51
Hess' Law ...........................................................................................................................57
Determination of Iron by Titration with Permanganate.....................................................63
Conductivity of Solutions ..................................................................................................71
Activity Series ...................................................................................................................77
Models................................................................................................................................87
Boyle's Law: Pressure-Volume Relationships in Gases ....................................................95
Molar Volume of Hydrogen...............................................................................................99
Appendix A: Statistical Treatment of Data......................................................................103
Appendix B: Equipment Hunt .........................................................................................105
Chemistry 125 Laboratory Manual –2004-2005
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Revision Dec 2003
Chemistry 125 Laboratory Manual –2004-2005
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Scientific Measurement
Goals:
Use the balance to make quantitative measurements
Use precision and accuracy in measurements
Use average deviation and percent relative average deviation with measurements
Read rulers, graduated cylinders, and burets correctly
Background:
In this experiment we will gain an understanding as to how to make measurements using a balance.
The measurements made will be used to explore the limitations of measurements and then how the
data obtained from measurements can be presented in a useful way. We will mass 5 pennies made
before 1983 and then 5 pennies made after 1983. We will then calculate the mean, the average
deviation and the relative average deviation of the two sets of pennies. If your instructor chooses, the
data from the entire class can be used and the standard deviation calculated.
In addition to using the balance, you will also learn how to make proper measurements using a ruler,
graduated cylinder, and buret. You will use these pieces of laboratory equipment throughout the
semester. Therefore, an early understanding of how to take proper measurements with this
laboratory equipment is a crucial skill in the laboratory.
The assessment of experimental error:
Every measurement involves some measurement error. Errors can be classified in two ways:
determinate error or indeterminate error. Determinate errors, also called systematic errors, are errors
in your methods, equipment/materials, personal judgments, and simple mistakes. Determinate errors
can be eliminated with practice and proper attention to detail. Indeterminate errors or random errors
cannot be controlled or eliminated. Thus, there will always be some degree of error in
measurements.
Accuracy:
The error in a measurement is the difference between the true value of the quantity measured and the
measured value. Error calculated this way is known as absolute error. The smaller the absolute
error, the closer the measured value is to the true value and the more accurate the result. Accuracy is
a measure of the correctness of the measurement. Frequently, we wish to compare several
measurement errors of different quantities. In these cases, it is more useful to use relative error.
Relative error is defined as the error divided by the accepted value and multiplied by 100. This is
usually known as percent error:
% error =
(experimental - accepted)
accepted
x 100
Sometimes we do not have or know an accepted value. In those cases, percent error cannot be
calculated. We do not have an accepted value for this experiment. Therefore, we will not be
calculating percent error.
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Precision:
The key to significance in experimental measurement is repetition. Only with repeated
measurements can an experimenter have some confidence in the significance of the measurements.
Only if a measured quantity can be reproduced repeatedly can the experimenter have that
confidence. Precision is a quantitative measure of the reproducibility of experimental measurements.
It is how well repeated measurements of the same quantity agree with one another. Precision is
frequently measured in terms of the average deviation, which is determined by the following
process:
1. From a series of measurements (three or more) determine the average value.
_
∑ xi
x = mean =
n
2. For each measurement determine its deviation from that average value.
average deviation =
_
∑| xi - x |
n
where, n= the number of data points and xi= the individual data points
Note: the summations in the formulae above go from i = 1 to n.
3. Determine the average of the deviations without regard to sign.
Example:
Four measurements of the concentration of an unknown acid by titration were made. The following
results were obtained: 0.1025 M, 0.1018 M, 0.1020 M, 0.1024 M. Compute the average value and
the average deviation.
Solution:
1. Compute the average value by summing the four measurements and dividing by four.
(0.1025 + 0.1018 + 0.1020 + 0.1024)/4 = 0.102175 rounded off to 0.1022
2. Deviation from the average value of 0.1022:
0.0003 (.1025-.1022), .0004 (.1018-.1022), .0002 (.1024-.1022), .0002 (.1024-.1022). Sum of the
deviation = 0.0011/4=.0003.
3. Since this deviation represents an uncertainty in the measurements, the molarity of the unknown
acid is not precisely 0.1022 M but ranges from 0.1019 M to 0.1025 M and should, therefore, be
reported with the average deviation included, that is 0.1022 ± 0.0003 M. Further, to make the
measurements of precision more useful, the average deviations are put on a percentage basis by
determining the relative average deviation (r.a.d.). (See Appendix A, pg. 103 for more
information). This is the average deviation divided by the average value and multiplied by 100%.
r.a.d. =
average deviation
mean
x 100
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In the above example, the relative average deviation is:
0.0003
0.1022
x 100% = 0.3%
The precision of an experiment varies with the technique and/or the apparatus used. A number of
variables built into the method or design of the experiment can affect its precision.
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Procedures:
Part I. Pennies Mass Measurement
Pennies minted in the United States in 1981 and earlier years are significantly different from pennies
minted in 1983 or later. Is it possible to tell the difference between these two kinds of pennies by
weighing them?
1. Put on your lab coat and your goggles.
2. Work in pairs.
3. One student will receive 5 pennies minted 1981 or before. The other student will receive 5
pennies minted 1983 or later.
4. Receive instructions from your instructor for the use of the pan balances in the laboratory.
5. Weigh and record the mass of each of the 5 pennies received.
6. Record the mass and the dates of the 5 pennies of your partner.
7. In your report calculate the mean of each set of pennies.
8. In your report calculate the average deviation
9. In your report calculate the relative average deviation
Part II. Proper Measurements Using Lab Equipment
WORK INDIVIDUALLY FOR THIS PART OF THE LAB!!
Ruler
1. Obtain a ruler and examine the units and markings. Your ruler should look something like this:
Suppose you wanted to measure the bold line above this ruler. Clearly the line is longer than 5
cm and shorter than 6 cm. Furthermore, the ruler is calibrated with lines each representing 0.1
cm. Counting the lines reveals the line to be somewhere between 5.8 cm and 5.9 cm. You must
estimate between the lines. Always read one more digit than is marked.
Final reading = 5.88 cm (The underlined “8” is the estimated digit.)
2. Measure the items listed under the Ruler section on the data page for this lab. Be sure to include
proper units and make all readings to the correct number of decimals.
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Graduated Cylinders
3. Obtain a 10 mL and 100 mL graduated cylinder. Examine the units and markings. Your
graduated cylinders should look something like this (note: they are not drawn to scale):
Suppose you want to measure the liquid shown in the 10 mL graduated cylinder. Notice the
liquid is curved down. This is called the meniscus. Aqueous solutions will always curve down.
Mercury is a liquid that actually curves upward. For purposes of this lab and this course, you
will only be using aqueous solutions; therefore, always measure liquids from the bottom of the
meniscus. Clearly the line is between 3 mL and 4 mL. The 10 mL graduated cylinder is
calibrated with lines each representing 0.1 mL. Counting the lines reveals the line to be
somewhere between 3.5 mL and 3.6 mL. You must estimate between the lines. Always read
one more digit than is marked.
Final reading = 3.53 mL (The underlined “3” is the estimated digit.)
Reading the 100 mL graduated cylinder is accomplished in the same way; however, notice each
of the markings represents 1 mL rather than 0.1 mL. A quick look at the graduated cylinder
indicates the volume to be between 20 and 30 mL. Each line represents 1 mL, therefore, the
line is clearly between 26 mL and 27 mL. You must estimate between the lines. Always read
one more digit than is marked.
Final reading = 27.2 mL (The underlined “2” is the estimated digit.)
Notice, the number of decimal place you can read depends on how the lines are calibrated on a
particular piece of equipment.
4. Practice using the 10 mL and 100 mL graduated cylinders located in your desk. THEN read the
graduated cylinders in the back of the room to the correct number of decimal places.
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Buret (Pronounced “byur●ette”)
5. Burets will be used several times during the semester. The calibrations on a buret are the same
as a 10 mL graduated cylinder. Therefore, the buret will be read to the same number of decimal
places. However, the numbers on the buret are reversed. The zero is at the top and 50 mL is at
the bottom. Your instructor should comment briefly on the reason for this and the uses of the
buret. The reason for the reversed numbers will become clearer as you use the buret during the
semester. The buret should look like this:
The line is clearly between 2.9 and 3.0 mL. You must estimate between the lines. Always read
one more digit than is marked.
Final Reading = 2.97 mL (The underlined “7” is the estimated digit.)
6. Read the burets located at the back of the room.
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Scientific Measurements
Data and Calculations
Name _____________________________________
Sec _____
PART I.
Your data from the weighing of five pennies:
Date of the penny
1.
2.
3.
4.
5.
Average (mean) mass of five pennies
Average deviation:
Average mass ± average deviation:
Percent relative average deviation:
Mass of the penny (xi)
1.
2.
3.
4.
5.
Your partner's data from the weighing of five pennies:
Date of the penny
1.
2.
3.
4.
5.
Average (mean) mass of five pennies
Average deviation:
Average mass ± average deviation:
Percent relative average deviation:
Mass of the penny (xi)
1.
2.
3.
4.
5.
Sample Calculation:
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Part II. Proper Measurements Using Lab Equipment
Ruler – Measure the following items to the correct number of decimals using a ruler. Include
proper units.
1. The diameter of a 250 mL beaker = ________________
2. The diameter of a penny = _______________
3. The length and width of this piece of paper:
Length = _______________
Width = _______________
Graduated Cylinders – Measure the graduated cylinders located in the back of the room to the
correct number of decimals. Include proper units.
4. Volume of Graduated Cylinder A = _______________
5. Volume of Graduated Cylinder B = _______________
6. Volume of Graduated Cylinder C = _______________
7. Volume of Graduated Cylinder D = _______________
Buret – Measure the burets located in the back of the room to the correct number of decimals.
Include proper units.
8. Volume of Buret A = _______________
9. Volume of Buret B = _______________
10. Volume of Buret C = _______________
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Questions:
1. Does the data you collected during Part I. indicate that there is a significant difference between
the masses of pennies minted in 1981 and earlier versus those minted in 1983 and later years?
YES
or
NO
(Circle One Answer)
2. Fill in the appropriate data and complete the question below.
Pennies Dated 1981 and earlier
Average ± average deviation = _________________________
Pennies Dated 1983 and later
Average ± average deviation = _________________________
Comparing the averages and using the average deviations, support your answer from question
#1.
3. Below you will find 3 different rulers each with different calibrations. Measure the bold line as accurately
as possible, to the correct number of decimals, and with proper units.
Length = _______________
Length = _______________
Length = _______________
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Density Experiment
Goals:
Learn to use equipment (graduated cylinder, balance).
Read equipment properly (estimate between lines).
Learn to use significant figures.
Calculate % error.
⇒ Prelab Report on page 19. Turn in when you come to class!
Background:
In your study of chemistry, you will often have to make quantitative measurements and carry out
calculations based on your data. If a correct or accepted result is known, you can compare your
results to this correct result. This lab is an introduction to the techniques and mathematics you will
need to use throughout the course.
Significant Figures:
The following rules can be used to determine the appropriate number of significant digits in a
measurement or of a number as the result of a calculation. The significant figures in a measured
quantity include all of the certain digits plus the last digit is considered to be the uncertain digit (not
known precisely, guessed). The uncertainty in the last digit in a number is assumed to be ±1 unless
otherwise noted.
For example, suppose you measured the mass of a penny and reported it to be 2.496 g. The first three
digits are assumed to be exactly correct, but the last digit is assumed to have an uncertainty of at
least ±1. In other words the mass could be as low as 2.495 g or as high as 2.497 g. So the value of
the measurement is reported as (2.496±0.001) g or simply 2.496 g.
Often, further calculations must be done after collecting the data. Significant figures should also be
used properly in these calculations. Your final answer should reflect the accuracy of the original
measurement, rather than the number of digits displayed by the calculator. In general, the number of
significant figures in the least accurate measurement limits the number of significant figures in the
answer.
Density:
The density of a substance is defined as the mass per unit volume: D =
Mass
Volume
The units of density are commonly g/mL or g/cm3 for most solid and liquid substances (a milliliter
is the same as a cubic centimeter). Density is a physical property, which has a characteristic value
for pure substance.
Some representative values are:
Lead
11.3 g/mL
Silver
10.5 g/mL
Aluminum
2.70 g/mL
Ice
0.92 g/mL
Ethanol
Water (4o C)
Mercury
Cork
Chemistry 125 Laboratory Manual –2004-2005
0.79 g/mL
1.00 g/mL
13.6 g/mL
0.26 g/mL
Page 13
Measuring the density of an unknown substance can help you determine what the substance is. For
example, if an unknown piece of shiny gray metal were found to have a density of 2.70 g/mL, you
would strongly suspect that the metal is aluminum.
Percent Error:
When a correct or accepted value is known, you can determine if your measurements give an answer
in close agreement with the accepted value. A common way to do this is to report the percent error.
Refer to appendix A for discussion of percent error.
Procedures:
To determine density, mass and volume must be measured. Mass can be determined directly from a
balance. Volume can be measured directly (measuring dimensions and using geometric formulas) or
indirectly (water displacement).
Taring the Balance:
When weighing several small objects or a powder, it is impossible to place it directly on the balance
pan. It is easiest to place the objects in a container, and have the balance subtract out the weight of
the container. This process is called “taring.”
Place the empty container on the balance and press the bar on the front. The display should now
read zero (there will be some small fluctuations due to air currents). Remove the container from the
balance (the balance will read a negative mass after the container is removed) and place the material
to be weighed in the container. It is better to do this away from the balance to protect the balance
from spills. The container with the material should then be placed back on the balance. The readout
will give the weight of the material alone (the weight of the container has been subtracted out by the
balance).
Volume by Water Displacement:
A common technique for determining volume is by using water displacement. This is often used
when the volume of the solid is difficult to calculate from geometry or difficult to measure, such as
an irregular shape. The technique involves measuring the amount of water in a container (graduated
cylinder), adding the solid and remeasuring the water level. The water level will increase due to the
presence of the solid. The difference between the levels (before and after adding the solid) is the
volume of the solid.
A. Regular Solid - Direct Measurement
1. Obtain a regularly shaped (cylinder or cube) piece of metal.
2. Determine the mass of the metal piece using the balance (record all the digits from the balance,
do not round off).
3. Using a ruler, measure the dimensions (height, diameter or length, width, height). Estimate
between the lines on the ruler (if each line is 0.1 cm, estimate to the nearest 0.01 cm). Record
the measurements.
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4. Show your calculations for the volume of the metal piece. Remember to use significant figures
properly in recording your answer and give the appropriate units.
5. Show your calculations for the density of your metal piece. Remember to use significant figures
properly in recording your answer and give the appropriate units.
6. When you are finished, return the metal piece and ruler to the appropriate place on the cart.
7. Using the CRC handbook available in the lab, look up the densities of several metals and alloys.
Choose the one which seems to fit the appearance and density of your sample the best and
calculate the % error.
B. Regular Solid - Volume by Water Displacement
1. Obtain a cylinder made of metal.
2. Determine the mass of the metal cylinder (record all the digits from the balance, do not round
off).
3. Pour some tap water into your 100 mL graduated cylinder. The amount should be enough so that
it would completely cover the solid if it were submerged.
4. Record the level of the water. Estimate between the lines on the graduated cylinder - if the lines
represent 1 mL, you should estimate to 0.1 mL.
5. Tip the cylinder and gently slide the metal piece down the side (if it is dropped in, the glass will
crack). Record the level of the water with the metal included.
6. The volume of the metal is the difference between your two measurements. Show your
calculations for the volume of the metal. Remember to use significant figures properly in
recording your answer and give the appropriate units.
7. Show your calculations for the density of your metal piece. Remember to use significant figures
properly in recording your answer and give the appropriate units.
8. When you are finished, dry off the metal piece and return it to the appropriate place on the cart.
9. Using the CRC handbook available in the lab, look up the densities of several metals and alloys.
Choose the one which seems to fit the appearance and density of your sample the best and
calculate the % error.
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C. Irregular Solid
1. Obtain some marble chips (about 20-25 chips).
2. Determine the total mass of your marble chips.
3. Place an empty plastic weigh boat on the balance pan and press the bar on the front of the
balance to tare it. Remove the weigh boat and place the marble chips in it. Place the weigh boat
with the marble chips back on the balance. The readout should give the mass of the marble chips
alone.
4. Record the mass (record all the digits from the balance, do not round off).
5. The volume will be determined by the technique of water displacement (see above).
Pour some tap water into your 100 mL graduated cylinder. The amount should be enough so that
it would completely cover the solid if it were submerged.
6. Record the level of the water. Estimate between the lines on the graduated cylinder - if the lines
represent 1 mL, you should estimate to 0.1 mL.
7. Gently drop the marble chips into the graduated cylinder. Make sure there are no air bubbles
between the chips. Record the level of the water with the marble included.
8. The volume of the marble is the difference between your two measurements. Show your
calculations for the volume of the marble chips. Remember to use significant figures properly in
recording your answer and give the appropriate units.
9. Show your calculations for the density of marble. Remember to use significant figures properly
in recording your answer and give the appropriate units.
10. When you are finished, return the marble chips to the appropriate container on the cart. (They
are not placed back in the original container with the unused chips).
11. Using the CRC handbook, look up the density of marble. Calculate the % error.
D. Liquid
1. Obtain a small amount (about 10 mL or less) of salt water in a beaker.
2. Clearly, a liquid cannot be poured onto the balance in order to be weighed. It must be in a
container. The technique of taring the balance (see above) will be used to determine the weight
of the liquid without the container.
3. Place an empty 100 mL graduated cylinder on the balance pan. Press the bar on the front of the
balance. Remove the graduated cylinder and pour between 5 and 10 mL of the liquid in it (do
this away from the balance so that any spill will not damage the balance). Place the graduated
cylinder with the liquid back on the balance. The readout should give the mass of the liquid
alone.
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4. Record the mass (record all the digits from the balance, do not round off). Note that as the water
evaporates, the mass will decrease somewhat.
5. The volume can be read directly from the graduated cylinder. Record the level of the liquid.
Estimate between the lines on the graduated cylinder - if the lines represent 1 mL, you should
estimate to 0.1 mL.
6. Show your calculations for the density of salt water. Remember to use significant figures
properly in recording your answer and give the appropriate units.
7. In order to determine the density more precisely, the measurements will be repeated with a
different piece of equipment.
8. Place a clean, empty 10 mL graduated cylinder on the balance and tare. Pour between 5 and 10
mL of the salt water into the graduated cylinder (again, do this away from the balance). Place
the graduated cylinder with the liquid back on the balance. The readout should give the mass of
the liquid alone.
9. Record the mass (record all the digits from the balance, do not round off). Note that as the water
evaporates, the mass will decrease somewhat.
10. The volume can be read directly from the graduated cylinder. Record the level of the liquid.
Estimate between the lines on the graduated cylinder - if the lines represent 0.1 mL, you should
estimate to 0.01 mL.
11. Show your calculations for the density of salt water. Remember to use significant figures
properly in recording your answer and give the appropriate units.
12. When you are finished, pour the salt water down the drain. Rinse out your graduated cylinder
thoroughly so that the salt water does not dry inside and leave a residue.
13. Find out the accepted value of the density of the salt water (it may be written on the label of the
container or may be given by your instructor). Calculate the % error.
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Density Experiment
Prelab Exercise
Name __________________________
Sec ______
How many significant figures are in each of the following measurements?
1.005 mL _____
1.20 x 103 m
_____
0.012 g
0.007 mL
_____
15.00 cm _____
1.05 x 10-4 g
_____
10.00 g
0.00101 L
_____
_____
_____
1.000 x 106 cm _____
0.0510 L _____
Carry out the following calculations and round off the answer to the correct number of significant
figures:
1.009 g + 0.12 g =
_____________
10.00 g + 0.050 g =
_____________
5.2 mL + 10.0 mL =
_____________
23.00 g + 0.70 g =
_____________
5.012 g / 2.0 mL =
_____________
10.0 g / 3 mL =
_____________
15.072 g / 11.0 mL =
_____________
0.090 g / 0.1 mL =
_____________
(2.54 – 1.74) x 25.5 =
_____________
(2.91 – 2.70)/2.70 =
_____________
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Density Experiment
Data and Calculations
Name _______________________
Record measurements as accurately as you can.
Include units with all measurements and calculations.
Report all calculated answers to the correct number of significant figures.
Show work for calculations.
Sec _____
A. Regular Solid - Direct Measurement
Mass
_________________
Description of object ___________________
Dimensions: (Be sure to include the units)
Length
_________________
Width
_________________
Height
_________________
Diameter
_________________
Volume
________________
Density
________________
Show Calculations
Accepted value ________________
Percent error
________________
B. Regular Solid - Volume by Water Displacement
Mass _________________
Volume of water alone _________________
Volume of water and metal _________________
Volume of metal _________________
Density _______________
Reference Source for Accepted Value of Density: ____________________________
(Include the name, edition and page number)
Identification of Metal _______________
Accepted Value for Density _______________
Percent error ____________
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C. Irregular Solid
Show Calculations
Mass _________________
Volume of water alone _________________
Volume of water and marble chips_________________
Volume of marble _________________
Density of Marble _______________
Accepted Value for Density _______________
Percent error
____________
D. Liquid:
100 mL graduated cylinder
Mass of liquid _________________
Volume of salt water _________________
Density ____________
Percent Error _______________
10 mL graduated cylinder
Mass of liquid _________________
Volume of salt water _________________
Density ____________
Percent Error _______________
Accepted Value for Density ______________
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General Questions:
1. What is the limiting factor in determining the number of significant figures in your measurement of
density - mass or volume? If you wished to determine the density to more significant figures, would you
buy a better balance (to measure mass more precisely) or try to improve the measurement of volume (for
example, by using calipers instead of a ruler, or a more finely calibrated graduated cylinder)?
2. In determining the density of the liquid, you weighed the liquid in a graduated cylinder and then
determined the volume by reading the graduated cylinder. Suppose that you weighed the liquid in a
beaker, and then transferred it to a graduated cylinder in order to determine the volume. Which
approach would be more likely to give an accurate result? (Hint: What errors could be introduced by
transferring liquid from one container to another?)
3. Suppose, after you determine the density of your solid, you discover that it is hollow. How would this
affect your value for the measured volume (would your measured volume be higher or lower than the
actual)? How would this affect your value for the density (would the calculated value be higher or lower
than the actual value)?
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4. Suppose that some air bubbles are trapped in between the marble chips when you measure the volume
by water displacement. How would this affect your value for the volume (would your measured volume
be higher or lower than the actual)? How would this affect your value for the density (would the
calculated value be higher or lower than the actual value)?
5. Can water displacement be used to calculate the density of a cork? If so, how? If not, why not?
6. How does the density of 4 mL of water compare to the density of 2 mL of water? (Same, twice as
much, half as much?) Explain your answer.
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Separation of a Mixture
Goals:
Devise a scheme of separation based upon physical properties of substances.
Write up a laboratory report outlining the procedures and equipment used.
Cleanly separate a mixture and recover as much as possible of the components.
Background:
Physical properties of substances are those which can be measured without changing the identity of the
substance.
Examples of physical properties are density, color, melting point, magnetism and solubility. These
properties can be significantly different from one substance to another and can be made use of in order to
separate substances. For example, in doing laundry, clothes can be separated by color.
Chemical properties, which involve the change of the substance, are not as useful since the substance will
not be recovered in the same pure form. For example, in doing laundry, cotton materials are known to be
flammable (they react with oxygen). However, setting the clothes on fire will destroy them.
Materials and Procedures:
You will be given a mixture of five components. The mixture contains copper(II) sulfate, sand, iron filings,
foam pieces, and boric acid. The following physical properties are known about the substances:
copper(II) sulfate: soluble in both hot and cold water; blue solid; density: 2.3 g/mL
boric acid: soluble in hot but not cold water; white solid; density: 1.4 g/mL
sand: not soluble in water; brown, tan and white solid; density is greater than 1.0 g/mL
iron filings: not soluble in water; black solid; density: 7.9 g/mL
foam pieces: not soluble in water; white solid; density is less than 1.0 g/mL
You are to design an experimental procedure that will separate the five components into their natural states
(all should be recovered as solids).
You will be given a test tube containing one gram of the mixture. After separation, place the components
into small baggies labeled with your name and the name of the component.
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A flow chart is often helpful in clarifying the procedures to be used. For example:
Dirty Clothes (White, delicate and permanent press)
Examine each piece, separate by color
Dirty colored clothes
(Delicate and permanent press)
Dirty White Clothes
Place in washing machine
Add detergent
Add Bleach
Set machine to hot water wash
Run washing machine
Read label of each piece
Separate delicate from
permanent press
Clean white clothes
Dirty colored permanent
press clothes
Place in washing machine
Add detergent
Set machine to warm water wash
Run washing machine
Clean colored permanent
press clothes
Dirty colored delicate
clothes
Place in washing machine
Add special detergent
Set machine to delicate
Run washing machine
Clean colored delicate
clothes
The flowchart describes the procedures used at each step, and gives a clear diagram of the separation. At
each step, it is clear which component is being separated and which still remains in the mixture.
Chemistry 125 Laboratory Manual –2004-2005
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Separation of a Mixture
Name ____________________________
Sec _____
Flow Chart:
Design a flow chart of your separation. At this point, check with your instructor before proceeding.
Chemistry 125 Laboratory Manual –2004-2005
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Procedure:
List the steps in numerical order for your procedure. This should be a verbal explanation of your flow chart.
Describe the procedures you are using. Correct terms should be used for equipment.
The procedure should be written clearly enough that another person could understand and follow the
instructions.
Chemistry 125 Laboratory Manual –2004-2005
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Equipment:
List the equipment that you will need to accomplish the separation.
Conclusions:
1. Discuss the sources of error in your experimental design. How might you do things differently to
eliminate these errors?
2. List five specific examples from your experiment, which demonstrate how physical (not chemical)
properties can be used to separate a mixture.
Chemistry 125 Laboratory Manual –2004-2005
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3. Using the information given, indicate how a mixture of the compounds might be separated:
Use a flowchart for your answers.
PbCl2 and AgCl
PbCl2: white solid, density 5.85 g/mL, soluble in hot water but not in cold water
AgCl: white solid, density 5.56 g/mL, not soluble in hot or cold water
Ice, methanol (CH3OH), and Styrofoam
ice: density 0.9 g/mL, melting point 0°C
methanol: density 0.79 g/mL, boiling point 65°C, melting point –97.8 °C, soluble in water
Styrofoam: white solid, density 0.47 g/mL, not soluble in water
Fe2S3 and Fe(NO3)3
Fe2S3: green solid, density 4.3 g/mL, decomposes upon heating, not soluble in water
Fe(NO3)3: violet solid, density 1.68 g/mL, decomposes upon heating, soluble in water
Chemistry 125 Laboratory Manual –2004-2005
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Determination of a Chemical Formula
Goals:
Carry out a chemical reaction and isolate the product.
Determine the empirical formula of a compound.
⇒ Prelab Report on page 33. Turn in when you come to class!
Background:
A formula of a chemical compound is a description of what chemical elements are present in the compound.
When the elements are nonmetals, the compound is usually made up of molecules. The molecules are atoms
covalently bonded together, such as H2O, water, or C6H6, benzene.
When metals and nonmetals combine to make a compound, there are usually positive and negative ions
present. The formula is the simplest ratio of ions. There are no individual molecules in which a certain
cation is bonded to a certain anion. For example, in sodium chloride each sodium ion is surrounded by six
chloride ions and each chloride ion is surrounded by six sodium ions. The formula is NaCl because the
smallest whole number ratio is sodium and chloride is one -to-one. In calcium chloride the formula is CaCl2
because for each calcium ion there are two chloride ions present.
Because atoms are so small, we cannot discover the formula of each chemical substance by counting atoms
or ions. Instead, larger quantities of chemical substances are used. The masses of elements in chemical
compounds are found. The analysis can be done in two ways: (1) The pure compounds are broken down and
the amount of each element is found; (2) Elements are combined to give a compound and the amounts of the
elements that react are found. Using known molar masses, gram amounts of the elements are converted to
moles. The smallest whole number ratio of moles of each element is the chemical formula.
For example:
A compound contains 2.65 g aluminum and 2.35 g oxygen.
Convert grams to moles: 2.65 g Al (1 mol Al/27.0 g Al) = 0.09815 mol Al
2.35 g O (1 mol O/16.0 g O) = 0.1469 mol O
Find the mole ratio:
0.09815 mol Al/0.09815 = 1.000 Al
0.1469 mol O/0.09815 = 1.497 O
Convert, if necessary, to a whole number ratio:
1.000Al x 2 = 2 Al
1.497 O x 2 = 2.994 O
The formula for the compound is Al2O3.
Chemistry 125 Laboratory Manual –2004-2005
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Procedure:
You will react zinc with iodine to form zinc iodide. The mass of each reactant is measured, as well as the
mass of the product and any leftover reactants. By calculating the mole Zn : mole I ratio, you will
determine the chemical formula of zinc iodide.
Methanol is poisonous and flammable. Avoid skin contact and avoid breathing the vapors.
(Review the procedure for tarring a balance in the Density experiment).
1. Weigh out approximately 1.0 g of zinc metal.
2. Weigh out approximately 1.0 g of iodine solid. DO NOT LET ANY IODINE FALL ON THE
BALANCE! It reacts with the metal and damages the balance.
3. Weigh a clean dry 125 mL Erlenmeyer flask and record the mass.
4. Add the zinc to the flask, reweigh, and record the mass.
5. Add the iodine to the flask, reweigh, and record the mass.
6. Add approximately 25 mL of methanol to the flask. Cover the flask with foil. Place the flask on a hot
plate in the fume hood and adjust the temperature to heat the methanol to boiling. (Try about 3 on the
dial of the hot plate for this temperature.) Swirl the flask occasionally. The reaction should be complete
in approximately 15-20 minutes. There will still be some particles of zinc metal left in the flask, but the
iodine should have disappeared.
7. Weigh a clean 250 mL beaker and record the mass.
8. Carefully pour the liquid from the reaction flask into the beaker. Do not let any zinc leave the flask and
fall into the beaker. To make certain that all the zinc iodide is transferred to the beaker, add 5 mL of
methanol to the zinc left in the flask and swirl. Pour the liquid into the beaker with the zinc iodide liquid
- again, do not let any particles of zinc leave the flask. Repeat the washing with another 5 mL of
methanol.
9. Place the flask with the unreacted zinc back on the hot plate and allow it to dry. When it is dry, let the
flask cool and then weigh the flask with the remaining zinc metal.
10. Place the beaker with the zinc iodide and methanol liquid on the hot plate. Adjust the temperature so
that the methanol evaporates, leaving the zinc iodide in the beaker. Make sure all the methanol
evaporates and the product is dry. Allow the beaker to cool and then weigh the beaker with the zinc
iodide product.
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Determination of a Chemical Formula
Pre-lab Exercise
Name ___________________________
Sec _____
Lead forms two different compounds with chloride - one is an oily yellow liquid (compound A), the other is
a white solid (compound B). A chemist adds a sample of compound A to a test tube with a mass of 25.042
g (empty weight). On re-weighing the test tube with compound A he found the mass to be 25.529 g. On
heating, he found that some chlorine gas escaped and compound A was converted to compound B. He
found that the mass of the test tube and compound B was 25.429 g. Compound B was then heated in the
presence of hydrogen gas, reducing the lead to the metallic elemental form and driving off the chloride as
HCl gas. The weight of the remaining lead and the test tube was found to be 25.331 g.
What is the mass of compound A?_____________________
What is the mass of compound B?_____________________
What is the mass of the lead?_____________________
The lead left in the test tube at the end was present in both compounds A and B, so they both contain the
same amount of lead. The rest of the mass must be due to chlorine.
What is the mass of chlorine in compound A?_____________________
What is the mass of chlorine in compound B?_____________________
The mole to mole ratio will give the simplest formula of the compounds.
How many moles of lead were in the compounds?_____________________
How many moles of chlorine were in compound A?_____________________
How many moles of chlorine were in compound B?_____________________
What is the formula of compound A?_____________________
What is the formula of compound B?_____________________
Show your calculations:
Chemistry 125 Laboratory Manual –2004-2005
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Chemistry 125 Laboratory Manual –2004-2005
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Determination of a Chemical Formula
Data and Calculations
Name _______________________________
Before Reaction:
Sec _____
Show Calculations:
Mass of empty Erlenmeyer flask ______________________
Mass of flask and Zn
______________________
Mass of flask, Zn and I
______________________
After Reaction:
Mass of empty beaker
______________________
Mass of beaker and zinc iodide
______________________
Mass of flask with leftover Zn
______________________
Calculations:
Before reaction:
Mass of Zn
______________________
Mass of I
______________________
Moles of I
______________________
After Reaction:
Mass of zinc iodide
______________________
Mass of leftover Zn
______________________
Mass of Zn that reacted
______________________
Moles of Zn that reacted
______________________
Calculate the mole I reacted: mole Zn reacted ratio.
According to your data, what is the simplest formula for the compound zinc iodide?
Chemistry 125 Laboratory Manual –2004-2005
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A way to check your technique in this experiment is to see if mass is conserved.
Before Reaction:
Mass of Zn before reaction
______________________
Mass of I before reaction
______________________
Total mass before reaction
______________________
After Reaction:
Mass of zinc iodide
______________________
Mass of leftover Zn
______________________
Total mass after reaction
______________________
Questions:
1. In the procedures, why dissolve the iodine in methanol? Why not mix the two solids and allow them to
react?
2. Suppose some leftover zinc metal was transferred to the product. How would this affect your
calculations, and how would the formula of the product be affected? (Would the formula have a higher
I:Zn ratio or lower?)
3. Suppose a non-reactive impurity was present in the zinc at the start of the reaction, and remained in the
flask with the leftover zinc. How would this affect your calculations, and how would the formula of the
product be affected?
4. What changes did you observe that indicated a chemical reaction?
5. What physical properties of zinc metal and zinc iodide did you use in order to separate the mixture?
Chemistry 125 Laboratory Manual –2004-2005
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Solubility Rules and Writing Formulas
Goals:
Accurately record observations of the result of mixing two salt solutions.
Write net ionic equations for the reaction of salt solutions.
Identify an unknown by reaction with various salt solutions.
Background:
When a salt dissolves in water to form an aqueous solution, it exists in solution as individual ions. For
example, sodium phosphate (Na3PO4) does not exist as an aqueous solution of Na3PO4 molecules, but rather
as separate Na+ (sodium ions) and PO43¯ (phosphate ions) surrounded by water.
H2O
Na3PO4(s) ⎯→ 3Na+(aq) + PO43¯ (aq)
We can write the equation for the dissolving of silver nitrate as well:
H2O
AgNO3(s) ⎯→ Ag+(aq) + NO3¯ (aq)
Na +
Na +
Na +
3PO 4
PO 43Na + Na +
Na + PO 3-
NO 3 Ag +
Ag+ NO 3 NO 3NO 3-
Ag +
Ag+
4
If two solutions, each containing a soluble salt, are mixed, then the ions
from one solution are free to mix with the ions from the other. When
silver nitrate and sodium phosphate are mixed, the species present before a reaction takes place are:
Na+(aq) + PO43¯ (aq) + Ag+(aq) + NO3¯ (aq)
We can mix a solution of aqueous sodium phosphate and aqueous silver nitrate. The products are aqueous
sodium nitrate and solid silver phosphate. We get the complete balanced equation below:
Na3PO4(aq) + 3AgNO3(aq) → Ag3PO4(s) + 3NaNO3(aq)
Rewriting the equation using the ionic species in the solution before and after they are
mixed, we get the equation below. Silver phosphate is not a soluble salt and will
precipitate (the solid that forms) from the solution. The aqueous solution that remains
contains sodium ions and nitrate ions. This equation is sometimes called the complete
ionic equation.
NO3Na+
Na+ NO3Ag3PO4
Ag3PO4
3Na+(aq) + PO43¯ (aq) + 3Ag+(aq) + 3NO3¯ (aq) → 3Na+(aq) + 3NO3¯ (aq) + Ag3PO4(s)
If the resulting salts are soluble in the products side of the equation (such as sodium nitrate above) then no
reaction has really taken place between these ions and they are the spectator ions. Thus the product of this
chemical reaction is silver phosphate. The solution that remains at the end of the reaction contains the
spectator ions Na+ and NO3¯. Because these spectators do not participate in the chemical reaction that has
taken place, we do not need to write them in the overall reaction. Since the reactants are present as aqueous
ions and the product is a solid, a reaction is said to occur. We can cancel the spectator ions out of the
equation and rewrite the chemical reaction as the net ionic equation as shown below.
3Ag+(aq) + PO43¯ (aq) → Ag3PO4(s)
Chemistry 125 Laboratory Manual –2004-2005
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The insoluble silver phosphate is seen as a fine yellowish solid, which will slowly settle to the bottom of the
container.
One method of establishing the solubility of salts is to mix solutions of known composition and observe the
resulting mixture. If a precipitate forms, then at least one of the salts produced is insoluble. We could say,
based upon our example, that silver phosphate is insoluble. But does one of the two ions have more
influence on the reaction than the other? Are all silver salts soluble except the phosphate? Or are all
phosphates soluble except the silver? In this experiment we will examine the results obtained when nine
anions (in solution as sodium salts) are mixed with nine cations (in solution as nitrate salts).
Procedure:
Obtain one sheet of acetate with a 10x10 grid and obtain one set of the nine anions and one set of the nine
cations. (There should be one set of each per table.) Each of the salts is present as 0.1 M aqueous solutions
in bottles with a screw top and dropper.
1. Put an acetate sheet over a piece of white paper, and then put one drop of each of the anions (NaCl, NaI,
etc.) in each of the squares of the proper column on the acetate sheet.
2. Place one drop of each of the cations (AgNO3, KNO3, etc.) in each square of its proper row on the
acetate sheet. Drop the solutions onto the sheet carefully. DO NOT touch the dropper to the sheet or
the solutions in the bottles will become contaminated. If you make a mistake, blot up the liquid with
a KimWipe.
3. Most of the drops will remain clear, indicating that the salts formed are soluble and no reactions
occurred; but several solutions change color and form a precipitate. Often it helps to see the precipitate
if the white sheet is carefully slipped from under the acetate. Note on your report sheet any color
changes and/or precipitate formation. There may be changes due to further reactions with vapors or
light as time passes. On your report sheet, record only the reactions that occur. Leave the squares
empty where no reaction occurs.
4. Write equations for each mixture which resulted in the formation of a precipitate on the report sheet.
5. Develop a set of statements or rules about the solubilities of ions based upon your observations (e.g., all
nitrates are soluble). Write these solubility rules on a separate sheet of paper. Group your observations
together so that you have no more than four solubility rules.
6. Obtain an unknown from your instructor. Record the unknown letter on your report. Your unknown
contains both an anion and a cation. Following the procedure you used previously, place a drop of the
unknown in each box in the column and the row labeled “unknown.” React the unknown with a drop of
each anion and cation in the appropriate square. Compare the reactions of the unknown with those of
the known anions and cations, and determine which column of known anions matches the column of the
unknown, and which row of cations matches the row of the unknown. Note that the entire row or
column must match, not just some of the squares. It may not be possible to completely identify the
unknown, but you should be able to narrow down the possibilities (e.g., your unknown may contain the
iodide ion as the anion and either the sodium or the ammonium ion as the cation. The cation cannot be
narrowed down any further based upon the reactions you have observed in this lab).
Chemistry 125 Laboratory Manual –2004-2005
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Solubility Rules
Name ______________________________
Sec _____
Solubility Table
anions
→
cations
↓
PO43¯
C2H3O2¯
NO3¯
Cl¯
SO42¯
I¯
OH¯
CO32¯
Unknown
anion
K+
Ag+
Ba2+
Na+
Ni2+
NH4+
Mg2+
Al3+
Zn2+
Unknown
cation
Chemistry 125 Laboratory Manual –2004-2005
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Data:
In the space below, write balanced net ionic equations for the formation of any precipitates you observed in
the experiment.
Unknown Letter _________________
Which cation was in the unknown? ________________
Which anion was in the unknown? _________________
Write the name and formula of the ionic compound in the unknown (e.g., if your unknown contained Ba2+
and I-, you would write Barium Iodide, BaI2).
Briefly describe how you reached a decision about your unknown. You may use equations, but don’t limit
your explanation to them.
Chemistry 125 Laboratory Manual –2004-2005
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Questions:
1. Suppose you were asked to prepare solid silver chloride from a solution of silver nitrate. Describe the
chemical reaction you would carry out (include any necessary reagents) and how you would separate the
solid product from the other materials.
2. List your solubility rules on this page. Write rules that you can draw from this laboratory experiment.
Do not simply restate the rules that you have learned in lecture. In other words, derive rules that are
based on your experimental results. Group your observations together so that you have no more than
four solubility rules.
Chemistry 125 Laboratory Manual –2004-2005
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Chemistry 125 Laboratory Manual –2004-2005
Page 42
Acid - Base Titration
Goals:
Learn to read a buret.
Learn how to accurately carry out a titration.
Calculate the molarity of an unknown acid or base from titration data.
⇒ Prelab Report on page 47. Turn in when you come to class!
Background:
In this experiment, you will be carrying out the following acid-base neutralization reaction:
HCl(aq) + NaOH(aq) → H2O(l) + NaCl(aq)
In this reaction, one mole of acid will react exactly with one mole of base. This reaction can be used to
determine the concentration of an unknown acid (or base) by reacting the unknown with a base (or acid) of
known concentration.
In this experiment you will titrate several different samples of HCl(aq) with NaOH(aq). Titration is a
laboratory technique where a carefully measured amount of one solution is added from a buret to a fixed
amount of another solution in an Erlenmeyer flask. In today's experiment, the buret is filled with NaOH(aq).
The HCl(aq) samples are placed in Erlenmeyer flasks along with a color indicator. Sodium hydroxide
solution is added to the Erlenmeyer flask until a color change occurs. This is called the endpoint.
The following information is required to determine the concentration of the unknown acid:
1. The concentration of the base - this is determined through a careful titration with a precisely measured
amount of acid (a process called standardization). You will be using a base, which has been previously
standardized so that the molarity of the base is accurately known.
2. The volume of base used in the reaction - you will use a buret to measure the amount of base used in
the reaction. Knowing the molarity of the base and the volume of base used, it is possible to calculate
the number of moles of base used in the reaction. According to the titration reaction shown above, the
number of moles of base added at the endpoint is equal to the number of moles present in the sample
initially.
3. The volume of acid used in the reaction - you will use a buret to measure an accurate volume of the
unknown acid. The molarity of the acid solution can then be calculated from the moles of acid
(determined in the titration) and the volume of acid.
4. A means of determining the endpoint (the point at which just enough base has been added to
neutralize the acid) - a substance called an indicator will be added to the reaction mixture. An indicator
is a molecule which changes color as the concentration of acid (H+) changes. The indicator we will use
is called phenolphthalein. It is colorless in acid solutions, and pink in alkaline solutions. As NaOH is
added to your HCl sample, hydrogen ions from the acid are removed from solution, and the indicator
will change color (becoming pink). The endpoint is reached when the indicator just changes color. Near
the endpoint the color change may occur very suddenly when a single drop of base is added to the flask.
Chemistry 125 Laboratory Manual –2004-2005
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Procedure:
Each student will work individually on this experiment.
A standardized solution of sodium hydroxide will be used to titrate a hydrochloric acid solution of
unknown concentration. You will do four titrations. Analysis of your titration data will allow you to
determine the molar concentration of HCl in the solution. The relative average deviation (see Appendix
A, "Statistical Treatment of Data") will also be calculated to give a measure of the precision of your
results.
1. Preparation of Buret:
a. Put on your lab coat and goggles.
Obtain approximately 125 mL of standardized NaOH(aq) in a clean, dry 250 mL beaker. Record the
concentration of NaOH on your report sheet.
b. The burets have been filled with water and are inverted in a beaker of water. Remove the buret and
drain all of the water out.
c. Rinse the buret with 3 separate 5-mL portions of NaOH(aq). Make sure you tip the buret so that the
NaOH solution touches the entire inner surface of the buret. Let some of the NaOH solution flow
through the tip of the buret to make sure no water is left in the tip.
d. Clamp the buret to a ring stand with a buret clamp.
e. Fill the buret with the standardized NaOH solution making certain there are no air bubbles in the tip
of the buret. Read the level of the solution in the buret (read the bottom of the meniscus). Be sure to
estimate between the lines to the nearest 0.01mL.
2. Preparation of HCl samples: You will prepare 4 HCl samples for titration. If you have more than one
Erlenmeyer flask, label each flask. Each sample should have a slightly different volume in the range of
25 - 30 mL. Accurately record (to the nearest 0.01 mL) the volume of each sample.
a. Clean a 250 mL Erlenmeyer flask and rinse it with deionized water. (Note that the flask does not
need to be dried after rinsing).
b. The HCl solution will be dispensed using the burets set up in the back of the room.
c. Make certain there are no air bubbles in the tip of the buret, then record the initial volume of HCl in
the buret. Deliver approximately 25 - 30 mL of HCl into your 250 mL Erlenmeyer flask. Record the
final buret reading.
d. Add 2-3 drops of phenolphthalein indicator to your Erlenmeyer flask. Place the flask containing
the HCl(aq) under the buret. Place a piece of white paper under the flask so that you can see the
color change more clearly once you begin the titration.
3. Titration:
a. Make sure you have recorded the initial volume of NaOH(aq) from the buret.
b. Begin adding NaOH(aq) to the Erlenmeyer flask while swirling the flask. Initially as NaOH is added
to the acid solution, you may see a temporary color change (colorless to pink). The endpoint of the
titration occurs when the indicator permanently changes color from colorless to a faint pink.
c. When you get close to the endpoint, add the NaOH one drop at a time. When you see the permanent
color change, stop adding NaOH. Use your deionized water bottle to rinse the tip of the buret into
the flask and rinse down the walls of the flask with a small amount of water.
d. Record the final volume of NaOH in the buret (to the nearest 0.01 mL).
e. The solution in the flask should be neutral and can then be poured down the sink.
4. Trials 2-4: Repeat the titration three more times as described above. Before beginning the second
titration, you may need to add more NaOH to the buret to make sure that you have enough to complete
the titration.
Chemistry 125 Laboratory Manual –2004-2005
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5. Clean up:
a. At the end of the experiment, carefully pour any NaOH (aq) remaining in the buret into the sink and
run tap water down the sink for one minute (dilute NaOH(aq) does not need to be neutralized prior to
disposal).
b. Thoroughly rinse the buret with water.
c. Fill the buret with deionized water (including the tip), put your finger over the end of the buret and
invert the buret into a 400 mL beaker that is half-filled with water. After you place the open-end of
the buret into the beaker of water, remove your finger. Make sure that the buret is securely clamped
to the ring stand.
Chemistry 125 Laboratory Manual –2004-2005
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Chemistry 125 Laboratory Manual –2004-2005
Page 46
Acid-Base Titration
Pre-lab Calculation
Name ___________________________
Sec ___
Turn in at the beginning of your laboratory session.
1. Oxalic acid is a solid that is often found as a hydrate, H2C2O4• XH2O where X is an integer. To
determine the value of X, a titration of a known amount of oxalic acid hydrate is carried out with a
strong base such as NaOH. The data from the titration can be used to calculate the amount of acid
present in the sample. The remainder of the mass is due to water.
5.012 g of oxalic acid hydrate is dissolved in enough water to make 100.0 mL of solution. This solution
is titrated with a standardized NaOH solution. It required 39.40 mL of 2.02 M NaOH to reach the
endpoint.
a. How many moles of NaOH were used in the titration?
Oxalic acid reacts with NaOH according to the following reaction:
H2C2O4(aq) + 2NaOH(aq)
→ Na2C2O4(aq) + 2 H2O(l)
b. Using your answer from part a and the balanced equation above, how many moles of oxalic acid
were neutralized in the titration?
c. The molar mass of oxalic acid is 90.04 g. How many grams of oxalic acid does your answer in part
b correspond to?
d. The remainder of the mass of the sample is due to water. How many grams of water are contained in
the sample of oxalic acid hydrate used in the titration?
e. How many moles of water does your answer in part d correspond to?
f. Calculate the ratio, moles of water to moles of oxalic acid. What is the formula of oxalic acid
hydrate? (i.e. what is the value of X in H2C2O4• XH2O?)
Chemistry 125 Laboratory Manual –2004-2005
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Chemistry 125 Laboratory Manual –2004-2005
Page 48
Acid-Base Titration Report
Name ____________________________________
Sec ______
Unknown Identification _______________
Molarity of standardized NaOH solution _______________
Trial 1
Trial 2
Trial 3
Trial 4
Final buret
reading NaOH
Initial buret
reading (NaOH)
Volume of NaOH
Moles NaOH
Moles Acid
Final buret
reading (acid)
Initial buret
reading (acid)
Volume of acid
Molarity Acid
Average Molarity of acid ___________________
Percent Relative Average Deviation (%RAD) _____________
Show all your calculations for one of the trials. Be sure to include units.
Chemistry 125 Laboratory Manual –2004-2005
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Questions:
1. State whether the calculated molarity will be higher, lower, or stay the same. Explain your answer.
a. The buret tip is not filled with NaOH solution before beginning the titration.
b. The acid contained impurities, which also reacted with the NaOH.
c. During the titration, some drops of NaOH stuck to the sides of the flask and did not react with the
acid.
d. The color at the end point is dark pink.
2. Potassium hydrogen phthalate (KHC8H4O4, molar mass = 204.23 g/mol ) is a monoprotic acid used to
standardize NaOH solutions. If you titrate 1.5556 g of this compound with 37.45 mL of NaOH solution,
what is the molarity of the NaOH solution? The reaction is:
KHC8H4O4 + NaOH → H2O + KNaC8H4O4
3. How many milliliters of 0.1229 M NaOH are needed to completely neutralize 50.00 mL of 0.1012 M
H2SO4 solution? Hint: First write a balanced neutralization reaction.
Chemistry 125 Laboratory Manual –2004-2005
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Cycle of Copper Reactions
Goals:
Carry out chemical reactions and make observations
Isolate the products of a reaction
Classify chemical reactions according to type
Determine the percent recovery of a substance after a series of reactions
Background:
Although there are thousands of chemical reactions, there are some broad categories that can be used to
classify many of them.
One such category is redox or oxidation - reduction. The loss of electrons is referred to as oxidation, and
the gain of electrons is referred to as reduction. Some of the reactions you will do in this experiment fit
this category.
Another category is metathesis, where ions exchange places: XY + WZ → XZ + WY. Some reactions
in this experiment will fall into this category.
No classification system can cover all reactions which occur, so there will be reactions which do not fit
into these two categories.
The cycle of reactions in this experiment is:
H2O
I. CuSO4•5H2O (s) →
CuSO4 (aq)
II. CuSO4 (aq) + 2NaOH(aq) → Cu(OH)2(s) + Na2SO4 (aq)
III. Cu(OH)2(s) → CuO(s) + H2O(l)
IV. CuO(s) + H2SO4(aq) → CuSO4(aq) + H2O(l)
V. CuSO4(aq) + Zn(s) → ZnSO4(aq) + Cu(s)
In many chemical reactions, you will need to separate the desired product from unwanted side products
or remaining reactants. Some of the separation techniques you have used in previous experiments (such
as filtration). In this experiment, you will carry out a cycle of reactions with copper. If the reactions are
carefully carried out and the product carefully purified, you should end up with the same amount of
copper as you started with.
Procedure:
Heat about 200 mL of deionized water in a beaker on a hot plate. This will be used later.
H2O
I. CuSO4 • 5H2O (s)
→
CuSO4 (aq)
To make 25.00 mL of a 0.400 M CuSO4(aq) solution, first calculate the mass of the CuSO4•5H2O(s)
needed to make 25.00 mL of this solution. (Molar Mass CuSO4•5H2O = 249.6) Record the mass of
the CuSO4•5H2O weighed. Measure out the appropriate amount of the CuSO4•5H2O and place in a
25.00 mL volumetric flask. Add enough water to fill the flask about half full and dissolve the
CuSO4•5H2O. Then fill the flask to the 25.00 mL mark and mix well. Transfer this solution to a 250
mL beaker.
Chemistry 125 Laboratory Manual –2004-2005
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II. CuSO4 (aq) + 2NaOH(aq) → Cu(OH)2(s) + Na2SO4 (aq)
While stirring the solution with a stirring rod, add 30 mL of 3.0 M NaOH solution. Record your
observations.
III. Cu(OH)2(s) → CuO(s) + H2O(l)
Heat the solution from part II on a hot plate. Stir gently to prevent the solution from becoming
locally overheated and boiling over (that is a superheated solution).
Record your observations.
When the reaction is complete ( as seen by the color change), take the beaker back to your lab bench
and allow the solution to cool and settle. Carefully pour off the liquid from the solid. Add about 200
mL of hot deionized water, allow the solid to settle, and pour off the liquid. This washing will
remove other materials in solution, leaving the copper oxide.
IV. CuO(s) + H2SO4(aq) → CuSO4(aq) + H2O(l)
Add 10 mL of 6.0 M sulfuric acid and stir. Record your observations.
V. CuSO4(aq) + Zn(s) → ZnSO4(aq) + Cu(s)
Place the beaker in the hood. Measure out approximately 2 g of Zn metal. Add the Zn metal to the
beaker in the hood. Record your observations.
At this point you should have recovered the solid copper metal. The copper needs to be isolated from
any remaining impurities and then dried. It is difficult to physically separate the remaining unreacted
Zn metal from the copper metal. We will use a chemical means to separate them. We will react the
zinc metal and put it in solution as its ion and then separate this solution from the solid copper metal.
When the solution is colorless, if there appear to be any silver-colored grains of unreacted Zn, add
10 mL of 6.0 M HCl (do not add the HCl immediately after adding the Zn) and gently warm the
solution. You will need to break up the chunks of copper metal so that the unreacted Zn metal can
completely react. When you can no longer see bubbles of hydrogen gas forming, pour off the liquid.
Transfer the copper to a porcelain evaporating dish. Wash the copper product with 5 mL of
deionized water, allow it to settle and pour off (decant) the liquid. Repeat the washing and
decantation two more times.
To thoroughly dry the copper metal, first wash it with 5 mL of methanol and decant. (Caution:
Methanol is highly flammable). Place the waste methanol in the container provided in the fume
hood.
Place the evaporating dish with the copper metal on a hot plate. The copper should be gently heated
until it is dry. If it is heated strongly for a long time, it will react with the oxygen in the air and give a
black copper oxide product instead of the desired copper metal.
Weigh an empty, clean 100 mL beaker and record the mass. Transfer the copper product to this
beaker and re-weigh. Record the mass of the beaker + copper metal. Your product should be dry and
have a typical copper color.
Turn in your product to your instructor for grading.
Chemistry 125 Laboratory Manual –2004-2005
Page 52
Cycle of Copper Reactions
Data and Calculations
Name _______________________________
I.
Mass of CuSO4•5H2O
Moles of Cu in sample
Mass of Cu in sample
Sec _____
____________________
____________________
____________________
Show your work for the calculations:
II. CuSO4 (aq) + 2NaOH(aq) → Cu(OH)2(s) + Na2SO4 (aq)
Observations:
What is the limiting reagent for this reaction?
What are the spectator ions in the solution?
What remains in the solution after the reaction is complete?
III. Cu(OH)2(s) → CuO(s) + H2O(l)
Observations:
What is removed by washing with the hot water?
Chemistry 125 Laboratory Manual –2004-2005
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IV. CuO(s) + H2SO4(aq) → CuSO4(aq) + H2O(l)
Observations:
What is the limiting reagent for this reaction?
What are the spectator ions in the solution?
What remains in the solution after the reaction is complete?
V. CuSO4(aq) + Zn(s) → ZnSO4(aq) + Cu(s)
Observations:
What is in the solution after the reaction with zinc is complete?
This is considered a redox reaction. Which element is being oxidized and which is being reduced?
Explain
Mass of beaker + recovered copper metal:
Mass of empty beaker:
Mass of recovered copper
Percent recovery =
Mass of recovered product
mass of initial copper in CuSO4
_________________________
_________________________
_________________________
x 100
Chemistry 125 Laboratory Manual –2004-2005
Show your calculations:
Page 54
What effect would the following have on the percent recovery of copper (increase, decrease or no
effect)?
_____
I.
Some of the liquid is splashed out of the beaker after the solution was made.
_____
II.
Too much NaOH is added to the solution.
_____
III.
In decanting the hot water wash, some of the solid is lost.
_____
IV.
Excess sulfuric acid is added.
_____
V.
Not enough Zn metal is added.
_____
V.
Some Zn metal remains in the copper product.
_____
V.
The copper product is not completely dried after the washing with methanol.
_____
V.
The copper product is heated too long and some copper oxide is formed.
What would be the proper disposal or clean up of the following:
1. A small amount of NaOH is spilled on the lab bench.
2. A small amount of copper(II) sulfate solution is spilled on the lab bench.
3. You take too much sulfuric acid from the stock bottle and have some excess that you can not use.
4. You take too much methanol from the stock bottle and have some left over.
Chemistry 125 Laboratory Manual –2004-2005
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Chemistry 125 Laboratory Manual –2004-2005
Page 56
Hess’s Law
Goals:
Calculate the heat of reaction from experimental measurements.
Become familiar with the additive property of heats of reaction (Hess’s Law).
Background:
Hess’s Law states that the change in enthalpy for any chemical reaction is independent of the path, or
series of steps, by which the reaction is carried out. The path from reactants to products could involve
one step or many, but the enthalpy change is the same.
Thermochemical data for reactions can be treated algebraically. This means that if two reactions can be
added together to give a third reaction, the two enthalpies can also be added together to give the
enthalpy of the third reaction. ∆H1 = ∆H2 + ∆H3.
In this lab, you will calculate the enthalpy of formation of magnesium oxide: Mg + 1/2 O2 → MgO
The reaction occurs readily and is easy to observe, but it is difficult to directly measure the heat
involved. By measuring the enthalpy for a series of reactions, the enthalpy of formation of magnesium
oxide can be determined indirectly.
The heat of a reaction can be determined by
q reaction = -q experiment = - [m c ∆T + c' ∆T]
m is the total mass of the solution
c is the specific heat capacity of the solution, assumed to be 3.97 J/goC
∆T is the change in temperature Tf - Ti
c' is the calorimeter constant
The calorimeter constant takes into account heat absorbed by the calorimeter. Assume it is zero if it is
not known.
The enthalpy for the reaction is then given by the heat of reaction (in kJ) per mole: ∆H = qreaction / mol
Chemistry 125 Laboratory Manual –2004-2005
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Procedure:
A. Mg(s) + 2 HCl(aq) → MgCl2(aq) + H2(g)
Obtain a calorimeter made of two nested Styrofoam cups. Make sure that the cups are clean and dry.
Determine the mass of the calorimeter.
Measure out 20. mL of 2.0 M HCl and pour into the calorimeter.
Determine the mass of the calorimeter and acid.
Measure the temperature of the acid using the probe. Make sure that you measure the temperature in
degrees Celsius.
Obtain a piece of magnesium ribbon, which is about 0.2 g. Clean the surface of the magnesium
ribbon with steel wool. Determine the mass of the magnesium.
Add the magnesium to the acid in the calorimeter and place the lid on the calorimeter. Insert the
temperature probe through the hole in the lid so that it is immersed in the liquid. Gently swirl the
calorimeter and monitor the temperature. (The reaction gives off heat for a while and should be
allowed to react until the temperature begins to fall.)
Record the highest temperature reached.
Pour the solution from the calorimeter into a large beaker at your bench for neutralization and
disposal at the end of the experiment. Rinse and dry the calorimeter.
B. MgO(s) + 2 HCl(aq) → MgCl2(aq) + H2O(l)
Measure out 20. mL of 2.0 M HCl and pour into the calorimeter.
Determine the mass of the calorimeter and acid.
Weigh out approximately 0.35 g of magnesium oxide. Record the precise mass.
Measure the temperature of the acid.
Add the magnesium oxide all at once to the acid in the calorimeter. Place the lid on the calorimeter.
Insert the temperature probe through the hole in the lid so that it is immersed in the liquid. Gently
swirl the calorimeter and monitor the temperature. Record the highest temperature reached.
Pour the solution from the calorimeter into a large beaker at your bench for neutralization and
disposal at the end of the experiment. Rinse and dry the calorimeter.
Neutralize the waste in the beaker with bicarbonate and wash it down the drain with water.
Chemistry 125 Laboratory Manual –2004-2005
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Hess's Law Experiment
Data and Calculations
Name: ___________________________ Sec ___
Part A:
Mg(s) + 2 HCl(aq) → MgCl2(aq) + H2(g)
Mass of empty calorimeter:
____________________
Mass of calorimeter + HCl:
____________________
Mass of HCl:
____________________
Mass of Mg:
____________________
Total mass of solution:
____________________
Initial temperature:
____________________
Final temperature:
____________________
∆T:
____________________
q reaction:
____________________
show work for qreaction
Volume of HCl:
____________________
Molarity of HCl (from label): ____________________
Moles of HCl:
____________________
show work for Moles of HCl
Mass of Mg:
____________________
Moles of Mg:
____________________
show work for Moles of Mg
Limiting reactant:
____________________
Determine ∆H using q reaction and converting to kJ and dividing by the moles of the limiting reactant:
∆H (kJ / mol)
____________________
Chemistry 125 Laboratory Manual –2004-2005
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Part B:
MgO(s) + 2 HCl(aq) → MgCl2(aq) + H2O(l)
Mass of empty calorimeter:
____________________
Mass of calorimeter + HCl:
____________________
Mass of HCl:
____________________
Mass of MgO:
____________________
Total mass of solution:
____________________
Initial temperature:
____________________
Final temperature:
____________________
∆T:
____________________
qreaction :
____________________
Show work for qreaction
Volume of HCl:
____________________
Molarity of HCl (from label): ____________________
Moles of HCl:
____________________
show work for Moles HCl
Mass of MgO:
____________________
Moles of MgO:
____________________
show work for Moles MgO
Limiting reactant:
____________________
Determine ∆H using q reaction and converting to kJ and dividing by the moles of the limiting reactant:
∆H (kJ / mol)
____________________
Chemistry 125 Laboratory Manual –2004-2005
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Combine the equations for reaction A
Mg(s) + 2 HCl(aq) → MgCl2(aq) + H2(g)
∆H = _______________
and reaction B
MgO(s) + 2 HCl(aq) → MgCl2(aq) + H2O(l)
∆H = _______________
With the equation for the formation of water:
1/2 O2(g) + H2(g) → H2O(l)
∆H = -286 kJ/mol
to obtain the equation for the formation of magnesium oxide:
1/2 O2(g) + Mg(s) → MgO(s)
∆H = ?
Use Hess’s Law to calculate ∆H for the formation of magnesium oxide.
Show how the equations are combined and show your work for the calculation of ∆H.
Find the heat of formation of magnesium oxide from the table of thermochemical data in the appendix of
your text.
∆H(text) =
Calculate your percent error. Show your work.
Chemistry 125 Laboratory Manual –2004-2005
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Questions:
1. Determining the sign of ∆H for reactions:
a. A thermometer is often used to measure the temperature change for chemical reactions. Is the
thermometer part of the system or part of the surroundings?
b. When the temperature reading on the thermometer increases as a reaction proceeds, is energy
moving from the system to the surroundings or from the surroundings to the system?
From the point of view of the chemical reaction, is this an exothermic or an endothermic process?
What should the sign of ∆H be?
2. Were the reactions in this experiment endothermic or exothermic? Use your data as evidence for
your reasons.
Reaction A:
Endothermic or Exothermic
Reason
Reaction B:
Endothermic or Exothermic
Reason
3. Is the formation of magnesium oxide exothermic or endothermic? Explain how you decide.
4. Two assumptions were made in this experiment that afforded us simpler calculations. List those two
assumptions.
Chemistry 125 Laboratory Manual –2004-2005
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Determination of Iron By Titration with Permanganate
Goals:
Practice techniques for doing accurate titrations.
Quantitatively determine the amount of iron (II) in an unknown.
⇒ Prelab Report on page 65. Turn in when you come to class!
Background:
In this experiment, you will be carrying out a redox reaction:
5 Fe2+ + MnO4- + 8 H+ → 5 Fe3+ + Mn2+ + 4 H2O
and using the titration techniques you practiced in a previous lab (Acid-Base Titration) to accurately
determine the concentration of the permanganate solution, and then to accurately determine the amount
of iron (II) in an unknown.
Previously, you carried out an acid-base titration where an indicator was used to determine the endpoint.
In this experiment, a color change occurs in the reaction itself, which can be used to signal the endpoint.
The permanganate ion is deep purple in color, and the manganese (II) ion is colorless. The endpoint of
the titration is signaled by a change from the colorless Fe2+ solution to a faint pink color (due to
unreacted permanganate).
The first part of the experiment will involve standardizing (precisely determining the molarity) of the
permanganate solution. This is done by carrying out the titration with a known quantity of an iron (II)
salt, Fe(NH4)2(SO4)2 • 6H2O. Once the permanganate concentration is known precisely, the solution can
be used to determine the amount of iron in an unknown iron (II) salt in the second part of the
experiment.
The same chemical reaction is done in the second part of the experiment. Only this time a different iron
compound is used. It contains iron as Fe2+, but the anion is not known. After doing a titration, the
concentration and volume of permanganate solution will be used to find moles of permanganate that
were required to oxidize all of the Fe2+. Moles of iron can be found using the mole ratio in the equation
above. Then moles of iron are converted to grams of iron. The percent of iron is the ratio of the mass of
iron to total mass of the compound, with the fraction converted to a percent.
Chemistry 125 Laboratory Manual –2004-2005
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Procedure:
Part A:
Note: Work individually.
1. A concentrated solution of potassium permanganate (approximately 0.2 M) will be provided.
You will need to dilute this to 0.02 M, and you will need 300 mL of the dilute solution.
Calculate the amount of concentrated permanganate solution you will need to obtain (show work
on next page).
Place this volume of 0.2 M KMnO4 in a clean 500 mL Erlenmeyer flask and dilute it to 300 mL
with deionized water. Tightly stopper the flask, and mix well by shaking and swirling the flask.
2.
Clean a buret. Then rinse the buret three times with 5-mL portions of your KMnO4 solution. Fill
the buret with the solution, making certain that there are no air bubbles in the tip of the buret.
You will probably need to read the buret scale at the top rather than at the bottom of the
meniscus because the solution is so deeply colored. Remember to estimate between the lines
when reading the buret.
3.
Standardize the permanganate solution:
Weigh out about 1.00 - 1.10 g of the known iron (II) salt, Fe(NH4)2(SO4)2 • 6H2O. Record the
precise mass. Transfer all the salt to a clean flask. Use deionized water to rinse the weigh boat,
if necessary. Add deionized water to the flask so there are about 20 mL of water. Add 8 mL of
6M H2SO4. Titrate immediately. Record the buret reading at the endpoint. Neutralize the
sample before disposal.
4.
Calculate the moles of iron in your sample (the molar mass of the salt is 392.16). Calculate the
moles of permanganate used to react with it according to the stoichiometry of the balanced
equation. Use the calculated moles of KMnO4 and the volume of solution measured from the
buret to calculate the molarity of the permanganate solution. You will need three trials that
agree within 1%.
Part B:
1.
Choose an unknown iron salt and record the identification letter of the unknown. Weigh out
about 0.50 - 0.70 g of the unknown salt. Record the precise mass. Transfer all the salt to a clean
flask. Use deionized water to rinse the weigh boat, if necessary. Add deionized water to the
flask so there is about 20 mL of water. Add 8 mL of 6M H2SO4. Carry out the same titration
technique as was used to determine the molarity of the permanganate in part A. Using the
molarity of the permanganate from part A and the volume you just measured, calculate the moles
of permanganate used to titrate the unknown salt. Calculate the moles of iron according to the
stoichiometry of the balanced equation. Using the atomic weight of iron, calculate the grams of
iron that were present in your sample of the unknown salt. Then calculate the weight percent of
iron in your unknown. You will need two trials that agree within 1%.
Disposal:
Neutralize the titrants with NaHCO3. The solutions contain H2SO4!!
The 0.02 M KMnO4 solution may be poured down the drain with plenty of water.
Chemistry 125 Laboratory Manual –2004-2005
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Determination of Iron by Titration
Pre-Lab Exercise
Name ______________________________________
Sec _____
NOTE: You must show your work to get credit for your answers.
1.
(2 pts.) Suppose you are asked to prepare 300 mL of 0.020 M KMnO4 solution by dilution of a
0.20 M stock solution of KMnO4(aq). How many mL of 0.20 M KMnO4 solution are needed to
prepare 300 mL of 0.020 M KMnO4 solution? (Show how you arrived at your answer.)
2.
(8 pts.) In an acidic aqueous solution, Fe2+ ions are oxidized to Fe+3 ions by MnO4–:
5Fe2+(aq) + MnO4–(aq) + 8H+(aq) → 5Fe3+(aq) + Mn2+(aq) + 4H2O(l)
Suppose that 1.507 g of the salt FeCl2 is placed in a 250 mL Erlenmeyer flask to which 50 mL of
water and 10 mL of 6M H2SO4(aq) are added.
a. Why is sulfuric acid added to the Erlenmeyer flask?
b. How many moles of Fe2+ are present initially in the Erlenmeyer flask?
c. How many moles of MnO4– are required to reach the endpoint in this titration?
d. How many mL of 0.1277 M KMnO4 solution are required to reach the endpoint?
Chemistry 125 Laboratory Manual –2004-2005
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Chemistry 125 Laboratory Manual –2004-2005
Page 66
Determination of Iron by Titration
Data and Calculations
Name ________________________________________
Part A
Trial 1
Trial 2
Sec _____
Trial 3
Trial 4
Mass of iron salt
Moles of Fe2+
Moles of KMnO4
Buret reading
initial
Buret reading
final
Volume of KMnO4
Molarity of KMnO4
Calculate the average of three trials that agree within 1%:
Show all your calculations for one of the trials:
Calculate the RAD for your KMnO4 standardized solution.
Chemistry 125 Laboratory Manual –2004-2005
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Unknown Identification (circle one): Salt A
or
Salt B
Molarity of KMnO4 (average value from previous page): _____________________________
Part B
Trial 1
Trial 2
Trial 3
Trial 4
Mass of iron salt
Buret reading
initial
Buret reading
final
Volume of KMnO4
Moles of KMnO4
Moles of Fe2+
Grams of Fe2+
Weight % of Fe2+
in sample
Average of two trials that agree within 1%:
Show all of your calculations for one of the trials:
Chemistry 125 Laboratory Manual –2004-2005
Page 68
Questions:
1.
It is assumed that the unknown salt contains only iron in the +2 oxidation state. If some of the
iron is in the +3 oxidation state, how would this affect your results? Would the % that you
calculate be too high or too low?
2.
Why do you not need to accurately measure the volume of deionized water in which you dissolve
the iron salt?
3.
Some titrations are done using indicators, substances that change color when hydrogen ions are
added to or taken away from the indicator molecule. The color change indicates the endpoint of
the titration. Why isn’t an indicator used in this titration?
4.
Reactions that involve electron transfer happen at very different rates, depending on the
substances involved in the reaction and the solvent. Some happen as many as 105 times per
second, and some happen billions of times slower. In a titration, does it matter how fast the
reaction occurs? Explain your answer.
5. Suppose the concentration of permanganate used was 0.20 M, instead of 0.020 M. How would
this affect the volume of solution used? How would this affect the accuracy of your titration?
Chemistry 125 Laboratory Manual –2004-2005
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Chemistry 125 Laboratory Manual –2004-2005
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Conductivity of Solutions
Goals:
Measure conductivities.
Classify materials as non-, weak or strong electrolytes using conductivity data.
Explain conductivity in terms of the presence of molecules or ions.
Background:
In general, bonds between atoms in molecules can be classified as covalent (sharing electrons) or ionic
(transfer of electrons). When an ionic compound is dissolved in water, it will dissociate into the
component ions. For example, when NaCl dissolves in water, the solution contains Na+ and Cl– ions
rather than units of NaCl. A compound which has covalent bonds, such as C6H12O6 (glucose), will
continue to exist as a neutral molecule when dissolved in water, rather than breaking apart.
A substance which exists as ions in solution is capable of conducting electricity as a result of the ability
of the ions to move in the solution and is called an electrolyte. A substance which does not give ions in
solution is a non-electrolyte. Electrolytes are further divided into two categories: strong and weak. A
strong electrolyte is one which is completely dissociated in solution: NaCl in solution is composed
entirely of individual Na+ ions and Cl– ions. A weak electrolyte (such as HF) is partially dissociated into
ions, but the majority of the species present in solution are neutral molecules. A solution of HF in water
will contain mostly neutral HF molecules with a small amount (5% or less) of H+ and F– ions.
By measuring the conductivity of solutions, it is possible to determine whether a compound exists in
water solution as ions (strong electrolyte, high conductivity), mostly molecules and some ions (weak
electrolyte, low conductivity), or molecules (non-electrolyte, no conductivity). Determining if a
compound is a strong, weak or non-electrolyte, in turn, gives information about the nature of the
chemical bonds present in the compound.
For example, aqueous NaCl is a strong electrolyte and the species in solution are: H2O, Na+ and Cl–.
(The sodium and chloride ions are present in equal amounts.) Aqueous HF is a weak electrolyte and the
species present in solution are: H2O, HF, H+ and F–. (Most of the HF is present as neutral molecules
rather than ions.) Glucose is a non-electrolyte and the species present in solution are: H2O, C6H12O6.
(No ions are present.) Note that the solvent is listed as it is an important component of the solution and
may have an effect on the behavior of the solute.
Chemistry 125 Laboratory Manual –2004-2005
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Conductivity of Solutions
Procedure:
Use of the Conductivity Meter:
Probe: The probe is connected by a cable to the meter. Between measurements the probe can be placed
on its stand.
To make a measurement, put the solution to be tested into a test tube or small beaker. Liquid must be
deep enough to cover the holes near the bottom of the probe. Place the probe into the liquid so that the
liquid goes into the holes. Wait until the reading is “Stable”. Record the reading.
Rinse the probe well with DI water, rinsing through the holes. Touch gently to a Kimwipe to blot. Do
not wipe. Go on to the next measurement or put the probe on the stand.
Readings: Conductivity will be measured in units of Siemens per centimeter (S/cm). The meter’s
digital display will show a number and a unit, probably either microSiemens (µS/cm) or milliSiemens
(mS/cm).
Use the following ranges to classify substances:
Conductivity
Electrolyte Type
less than 200 µS/cm
non-electrolyte
200 – 2000 µS/cm
weak electrolyte
greater than 2000 µS/cm
strong electrolyte
Chemistry 125 Laboratory Manual –2004-2005
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Experimental:
I.
Using clean test tubes, obtain a small amount (about 10-20 mL) of the solutions listed below.
Place the probe in the solution so that it contacts the liquid for the lower 1-2 inches of the probe.
Record the value of the conductivity, making sure to note the units. Use the value of the
conductivity and the chart given above to decide if the compound is a strong, weak or nonelectrolyte.
Deionized water, ethanol(l) (CH3CH2OH), ethanol(aq) (CH3CH2OH), acetic acid(aq)
(HC2H3O2), sucrose(aq) (C12H22O11), NaCl(aq), 0.1 M HCl, 0.1 M NaOH(aq), 0.1 M
NH3(aq)
Waste disposal: Acids and bases should be neutralized with baking soda before pouring them
down the drain.
II.
Obtain a few marble chips (about 6-8). Place half of them in one test tube and add a small
amount (5 mL) of 6 M acetic acid. Place the remaining marble chips in another test tube and add
a small amount (5 mL) of 6 M hydrochloric acid. Note the relative reactivity of the marble chips
with the two acids.
Waste disposal: The solutions remaining after reaction are acidic and should be neutralized with
baking soda.
III.
Reactions:
You will need to use droppers for this section.
A. Mix 3 mL (about 60 drops) of 0.1 M acetic acid with 3 mL (about 60 drops) of 0.1 M
ammonia in a test tube. Measure and record the conductivity of the mixture. Compare the
conductivity of the mixture (after reaction) to the individual conductivities of the acetic acid
and ammonia solutions you measured in Part I.
B. Obtain 3 mL (about 60 drops) of 0.1 M H2SO4. Measure and record the conductivity. Obtain
3 mL (about 60 drops) of Ba(OH)2. Measure and record the conductivity.
To the test tube containing the H2SO4, add 1 drop of bromthymol blue indicator. You will
use the color change of this indicator to determine when the reaction is complete. Add the
Ba(OH)2 drop wise to the acid until the indicator changes from the yellow color to the blue
color. Measure and record the conductivity at the endpoint. Dispose of the product and any
remaining Ba(OH)2 solution in the container provided.
Chemistry 125 Laboratory Manual –2004-2005
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Chemistry 125 Laboratory Manual –2004-2005
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Conductivity Experiment:
Data and Calculations
Name ________________________
I.
Sec ____
In the following table, record the conductivity measured for each solution. The electrolyte type
should be strong, weak or non according to the table given previously. The species in solution
should be a listing (in order of decreasing concentration) of the components of the solution.
Solution
Conductivity
Electrolyte Type
Species in Solution
Deionized water
Ethanol(l)
(CH3CH2OH)
Ethanol (aq)
(CH3CH2OH)
0.1 M acetic acid (aq)
(HC2H3O2)
Sucrose(aq)
(C12H22O11)
0.1 M NaCl(aq)
0.1 M HCl(aq)
0.1 M NaOH(aq)
0.1 M NH3(aq)
Chemistry 125 Laboratory Manual –2004-2005
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II. Which acid reacted more rapidly? ______________________________________________
Which acid had the higher conductivity (from Part I)? _____________________________
Give an explanation of why conductivity and reactivity should be related.
III. A. Conductivity of acetic acid (from Part I) ______________________________________
Conductivity of ammonia (from Part I) __________________________________________
Conductivity of mixture ______________________________________________________
Did the conductivity increase or decrease after reaction? ____________________________
Write the reaction that occurred when the two solutions were mixed:
Based on your reaction, give an explanation of the results you observed for the conductivity
change.
III.
B. Conductivity of sulfuric acid
___________________________________________________
Conductivity of barium hydroxide ______________________________________________
Conductivity of mixture ______________________________________________________
Did the conductivity increase or decrease after reaction?
_____________________________
Write the reaction that occurred when the two solutions were mixed:
Based on your reaction, give an explanation of the results you observed for the conductivity
change.
Chemistry 125 Laboratory Manual –2004-2005
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Activity Series
Goals:
Balance redox equations.
Develop an activity series based on observations of reactions.
Predict whether a redox reaction will occur based on an activity series.
⇒ Prelab Report on page 81. Turn in when you come to class!
Background:
Many chemical reactions are classified as redox, or oxidation - reduction. Oxidation refers to the loss of
electrons by an element, and reduction is the gain of electrons. For example, in the reaction
Cu2+(aq) + Zn(s) → Zn2+(aq) + Cu(s)
the copper (II) ion is reduced to copper metal (the ion gains 2 electrons); and zinc is oxidized to zinc ion
(the neutral metal atom loses 2 electrons to form the ion).
We can break this reaction into two half reactions:
an oxidation reaction:
Zn(s) → Zn2+(aq) + 2eand a reduction reaction:
Cu2+(aq) + 2e- → Cu(s)
Another example:
2 Ag+(aq) + Cu(s) → 2 Ag(s) + Cu2+(aq)
The silver ion is reduced, the copper metal is oxidized.
Note that the equation is balanced so that the total charge on the reactants is equal to the total charge on
the products. Copper metal loses two electrons to form the ion, and each silver ion gains one electron;
thus it requires two silver ions to react with one copper atom.
Again, we can write this as the sum of two half reactions:
oxidation:
Cu(s) → Cu2+(aq) + 2e–
reduction
Ag+(aq) + e– → Ag(s)
When adding the reactions together to obtain the overall reaction, the reduction reaction is multiplied by
two. The number of electrons on the reactant side in the reduction reaction then becomes equal to the
number of electrons on the product side in the oxidation reaction, so that the electrons cancel when the
two half reactions are added together.
Comparing the two reactions given above, we note that copper ion is able to remove electrons from the
zinc metal, and that silver ion is able to remove electrons from copper metal. We can place these
reactions in order of decreasing ease of oxidation:
Zn(s) → Zn2+(aq) + 2eCu(s) → Cu2+(aq) + 2eAg(s) → Ag+(aq) + eAny metal at the top of the list can be oxidized by the ions of the elements below it. From the list, we
would predict the following reaction would occur:
Zn(s) + 2 Ag+(aq) → 2 Ag(s) + Zn2+(aq)
but not
Ag(s) + Zn2+(aq) → no reaction (zinc ion will not oxidize silver metal)
Chemistry 125 Laboratory Manual –2004-2005
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This list is referred to as an activity series. The element at the top of the list is the most easily oxidized
and the most reactive. In this experiment you will observe several combinations of metals and ions and
develop an activity series based upon your observations.
Procedure:
Part I
Using a spot plate, set up the following mixtures of solid metals and metal ion solutions. Use very small
pieces of each metal and add enough solution to cover each metal (but do not overfill the wells in the
spot plate). Make sure the reactants are mixed thoroughly. Some of the reactions will be slow, and you
need to wait for 15-20 minutes in order to determine whether a reaction has occurred. Many of the
metal ions form colored solutions, and a change in color of the solution is indicative of a reaction. It
may also be possible to see a change in the solid metal as well (one metal dissolves and a different color
metal appears). You should also write down your observations of the rate of reaction - a faster reaction
would indicate a more active metal.
1
2
3
4
5
6
Metal
Cu
Cu
Cu
Zn
Zn
Zn
Solution
Fe2+
Zn2+
Ag+
Cu2+
Ag+
Fe2+
Metal
Fe
Fe
Fe
Mg
Mg
Mg
Mg
7
8
9
10
11
12
13
Solution
Cu2+
Zn2+
Ag+
Fe2+
Zn2+
Ag+
Cu2+
Waste Disposal: Do not place any solids in the sink!!
Use forceps to remove the metals from the spot plate. Rinse the copper, zinc, and iron pieces in running
tap water then throw them in the trash. After rinsing the Mg pieces, place them in a waste container for
magnesium.
Part II
Many metals can also be oxidized by the hydrogen ions present in water or acid (the hydrogen ions are
reduced to hydrogen gas). The relative rate of reaction can be used to judge the activity of the metal the more active metals will react rapidly with cold water, whereas less active metals may react with hot
water or with acid, and the reaction may be slower. Polish each surface well.
A. Place about 5 mL of cold water in each of four test tubes. Carefully add a small amount of the
following metals to the test tube and record your observations. (Adding a little phenolphthalein to
the test tubes will indicate reactivity if a color change is observed.) You will need to wait up to 10
minutes for reactions to occur.
Test Tube Metal
1
Ca
2
Mg
Test Tube
3
4
Metal
Zn
Fe
Chemistry 125 Laboratory Manual –2004-2005
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B. Using a spot plate, set up the following mixtures of solid metals and 1 M HCl(aq). Use small pieces
of each metal and add enough 1 M HCl(aq) to cover the metal. Record your observations.
1
2
Metal
Cu
Mg
3
4
Metal
Zn
Fe
Waste Disposal: Do not place any solids in the sink!!
Pour the contents of the calcium test (part A) in a container labeled “calcium test.” Place any unreacted
Mg pieces in a waste container for magnesium. Rinse the copper, zinc, and iron pieces in running tap
water (use forceps to hold the metal pieces) then throw them in the trash.
Chemistry 125 Laboratory Manual –2004-2005
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Chemistry 125 Laboratory Manual –2004-2005
Page 80
Activity Series
Prelab Exercise
Name ______________________________________
Sec _____
An experiment to determine an activity series was carried out using lead, gold, and lithium. The
following results were obtained:
Li(s) + cold water: rapid reaction, metal dissolved and produced bubbles of hydrogen gas.
Au(s) + cold water: no reaction
Pb(s) + cold water: no reaction
Pb2+(aq) + Li(s): rapid reaction, light colored lithium metal dissolved and darker colored lead metal
appeared
Au3+(aq) + Pb(s): slow reaction, dark colored lead metal disappeared and yellow colored gold metal
appeared
Li+(aq) + Au(s): no reaction
Write the oxidation half reactions for Pb, Au, and Li:
Pb(s) →
Au(s) →
Li(s) →
Which metal is most reactive and easiest to oxidize?
Which metal is least reactive and most difficult to oxidize?
Write the half reactions in order from easiest to oxidize (at the top) to most difficult to oxidize:
1.
2.
3.
Based on your activity series, will this reaction occur?
Pb2+(aq) + Au(s) →
Explain.
Chemistry 125 Laboratory Manual –2004-2005
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Chemistry 125 Laboratory Manual –2004-2005
Page 82
Activity Series Experiment
Data and Calculations
Name ___________________________ Sec ___
Part I
Observations:
Cu2+
Zn2+
Fe2+
Ag+
Cu
Zn
Fe
Mg
Chemistry 125 Laboratory Manual –2004-2005
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Part II
Observations:
A1: Ca + water
A2: Mg + water
A3: Zn + water
A4: Fe + water
B1: Cu + acid
B2: Mg + acid
B3: Zn + acid
B4: Fe + acid
Write the half reactions for the oxidation of each metal (Ag , Cu, Fe, Mg, and Zn). Place the half
reactions in order of decreasing ease of oxidation. As in the example in the introduction, the metal at the
top of the list will be oxidized by the ions of the metals below it. The metal at the bottom of the list
cannot be oxidized by any of the metal ions above it.
Activity Series:
Based upon your observations of the reactions of metals (Ca, Cu, Fe, Mg, and Zn) with water and acid,
develop an activity series in order from most reactive to least reactive. Enter each of the metals as an
oxidation half-reaction. Be sure to place the hydrogen half-reaction (H2(g) → 2H+(aq) + 2e-) in the
appropriate position in your activity series.
Activity Series:
Chemistry 125 Laboratory Manual –2004-2005
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Questions:
1.
Complete and balance the following redox reactions:
Mg(s) + Fe2+(aq) →
Zn(s) + Cu2+(aq) →
Cu(s) + Ag+(aq) →
2.
Complete and balance the following redox reactions:
Mg(s) + H+(aq) →
Ca(s) + H2O(l) →
(Note: H2O behaves as a combination of H+ and OH-)
3.
Of the ions that you used in this experiment (Fe2+, Zn2+, Cu2+, and Ag+), which ion(s) would
oxidize Mg metal but not Cu metal?
4.
Underground pipes are often made of iron, which will eventually be corroded by the water
present in the ground. A piece of magnesium is often attached to the pipe in order to protect the
iron pipe from corrosion. Based on your activity series, would the magnesium or iron react more
rapidly with the groundwater?
Chemistry 125 Laboratory Manual –2004-2005
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Chemistry 125 Laboratory Manual –2004-2005
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Models
Goals:
To use balloon models to illustrate the shapes of structures according to the Valence Shell Electron Pair
Repulsion (VSEPR) Theory.
To practice pencil and paper calculation of bond angles and molecular structures.
To use model kits to make structures of given molecules and ions.
Background:
Most elements will gain, lose or share electrons in order to form an octet. (There are exceptions: Some
elements such as boron will form an incomplete octet and some such as phosphorus or sulfur will form
an expanded octet.) The four pairs of electrons in the octet will be arranged in a way that will minimize
electron pair-electron pair repulsion. This repulsion will also determine the way in which incomplete
octets or expanded octets will be arranged. Thus, the geometrical arrangement of electron pairs around
an atom can be determined from the Lewis dot structure.
Electron pairs can be involved in a bond with another atom or can be nonbonding (lone pairs). While
both are used to determine the arrangement of the charge clouds, the shape of the molecule is
determined by the arrangement of the atoms. The nonbonding or lone pairs are not visible, but their
presence does affect the shape. For example, the oxygen atom in water (H2O) has four pairs of electrons
- two bonding and two nonbonding. The geometric arrangement of the electron pairs (or charge clouds)
is tetrahedral, as this minimizes the electron pair repulsion. The shape of the molecule H–O–H is bent,
rather than linear, since the nonbonding electrons affect the angle between the atoms.
Chemistry 125 Laboratory Manual –2004-2005
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Procedure:
Part I – Shapes of Electron Charge Clouds on an Atom
You will need three long balloons or six round balloons. Blow them up so that they are not quite fully
inflated and tie the ends. If the long balloons are used, twist near the middle to make two lobes. Then
twist another balloon at the middle and twist it together with the first one, making four lobes. Do the
same with a third balloon to make six lobes that are joined in the center. If round balloons are used, tie
six together, making six lobes.
On your report page, sketch the shape the balloons make around their center. Label the angles between
the lobes. Give the name of the arrangement.
With a pin, pop one of the lobes. The balloon lobes should move to a new arrangement that keeps all
lobes as far apart as possible. If the balloons do not move easily because of friction between their
surfaces, help them a little. On your report page, sketch the shape the balloons make around their center,
label the angles and write the name for this arrangement.
Continue to pop one lobe at a time. Sketch the resulting arrangement of the balloons and the angles
between lobes. Name each arrangement.
Part II – Lewis Structures to VSEPR Shapes
First draw the Lewis dot structure for the given formula.
Count the total number of electron charge clouds around the central atom. A charge cloud is a group of
electrons and can be a lone pair, a single bond or a multiple bond. We are interested in the number of
charge clouds, not how many electrons there are in each.
Decide on the correct three dimensional geometry for the number of charge clouds.
If there are lone pairs of electrons on the central atom, decide on the correct molecular geometry.
Following is a summary of some possible combinations of bonding electrons and lone pairs. Note that
the number of charge clouds is the sum of the number of bonds and the number of lone pairs.
Number of
Charge
Clouds
2
3
3
4
4
4
5
Number of
Bonds
Number of
Lone Pairs
2
3
2
4
3
2
5
0
0
1
0
1
2
0
6
6
0
Chemistry 125 Laboratory Manual –2004-2005
Arrangement of
Charge Clouds
Linear
Trigonal planar
Trigonal planar
Tetrahedral
Tetrahedral
Tetrahedral
Trigonal
bipyramidal
Octahedral
Molecular
Geometry
linear
trigonal planar
bent or V-shaped
tetrahedral
trigonal pyramidal
bent or V-shaped
trigonal
bipyramidal
octahedral
Page 88
Part III – Models
Use the list of molecules and ions below. Add to the list any others required by your instructor.
Complete the analysis of each substance on the report page. Then use a model kit to make a model of
each particle. Some of the atoms needed are not listed in the color codes on the lid of the model kits.
For these, just use any pieces that have the right number of holes. Have your instructor check each
model.
ClO2–
PbCl4
NF3
NO3–
CS2
OF2
CH3C*H3
CH3C*O2H
(Consider the * carbon to be the central atom.)
Chemistry 125 Laboratory Manual –2004-2005
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Chemistry 125 Laboratory Manual –2004-2005
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Models Report
Name ___________________________ Sec_____
Part I
Sketch of Shape and Angles
Name
Six Lobes
Five Lobes
Four Lobes
Three Lobes
Two Lobes
Chemistry 125 Laboratory Manual –2004-2005
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Chemistry 125 Laboratory Manual –2004-2005
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Part II
Formula
Lewis Structure
Total #
Electron
Pairs
Around
Central
Atom
# Bond
Pairs
Around
Central
Atom
# Lone
Pairs
Around
Central
Atom
Electron
Pair
Geometry
Molecular
Geometry
Sketch
CO2
BCl3
SO2
BF4–
NH3
H2O
PCl5
SF6
Chemistry 125 Laboratory Manual –2004-2005
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Part III
Formula
Lewis Structure
Total #
Electron
Pairs
Around
Central
Atom
# Bond
Pairs
Around
Central
Atom
# Lone
Pairs
Around
Central
Atom
Electron
Pair
Geometry
Molecular
Geometry
Sketch
ClO2–
PbCl4
NF3
NO3–
CS2
OF2
CH3C*H3
CH3C*O2H
Chemistry 125 Laboratory Manual –2004-2005
Page 94
Boyle’s Law: Pressure-Volume Relationship in Gases
Name ______________________
Sec _____
The primary objective of this experiment is to determine the relationship between the pressure and
volume of a confined gas. The gas we use will be air, and it will be confined in a syringe connected to a
pressure sensor (see Figure 1). When the volume of the syringe is changed by moving the piston, a
change in the pressure exerted by the confined gas results. This pressure change will be monitored using
a pressure sensor interfaced to a computer. It is assumed that temperature will be constant throughout
the experiment. Pressure and volume data pairs will be collected during this experiment and then
analyzed. From the data and graph, you should be able to determine what kind of mathematical
relationship exists between the pressure and volume of the confined gas. Historically, this relationship
was first established by Robert Boyle in 1662 and has since been known as Boyle’s law.
Pressure Sensor
Figure 1
MATERIALS
Macintosh or IBM-compatible computer
Serial Box Interface or ULI
Logger Pro
Vernier Pressure Sensor
20-mL gas syringe
PROCEDURE
1. Prepare the Pressure Sensor and an air sample for data collection.
• Plug the Pressure Sensor into Port 1 of a Serial Box Interface or ULI that is connected to a
computer.
• Open the side arm of the pressure sensor valve to allow air to enter and exit. Open its side valve by
aligning the blue handle with the arm that leads to the pressure sensor as shown in Figure 2.
Figure 2
• Move the piston of the syringe until the front edge of the inside black ring (indicated by the arrow in
Figure 3 on the next page) is positioned at the 10.0 mL mark.
• Close the side arm of the pressure sensor valve by aligning the blue handle with the side arm (see
Figure 3).
Pressure Sensor
Figure 3
Chemistry 125 Laboratory Manual –2004-2005
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2. Prepare the computer for data collection by opening “Exp 06” from the Chemistry with Computers
experiment files of Logger Pro. The vertical axis has pressure scaled from 0 to 2.5 atm. The
horizontal axis has volume scaled from 0 to 20 mL.
3. Click
Collect
to begin data collection.
4. Collect the pressure vs. volume data. It is best for one person to take care of the gas syringe and for
another to operate the computer.
• Move the piston to position the front edge of the inside black ring (see Figure 3) at the 5.0-mL line
on the syringe. Hold the piston firmly in this position until the pressure value stabilizes.
• When the pressure reading has stabilized, click Keep . Type “5.0” in the edit box. Press the
ENTER key to keep this data pair. Note: You can choose to redo a point by pressing the ESC key
(after clicking Keep , but before entering a value).
5. Repeat the Step 4 procedure for volumes of 7.5, 10.0, 12.5, 15.0, 17.5, and 20.0 mL.
6. Click Stop when you have finished collecting data. In your data table, record the pressure and
volume data pairs displayed in the Table window (or, if directed by your instructor, print a copy of
the Table window).
7. Examine the graph of pressure vs. volume. Based on this graph, decide what kind of mathematical
relationship you think exists between these two variables, direct or inverse. To see if you made the
right choice:
• Click the Curve Fit button, .
• Choose Variable Power (y = Ax^n) from the list at the lower left. Enter the value of n in the
Degree/Exponent edit box that represents the relationship shown in the graph (e.g., type “1” if
direct, “-1” if inverse). Click Try Fit .
• A best-fit curve will be displayed on the graph. If you made the correct choice, the curve should
match up well with the points. If the curve does not match up well, try a different exponent and
click Try Fit again. When the curve has a good fit with the data points, then click OK .
8. Once you have confirmed that the graph represents either a direct or inverse relationship, print a
copy of the Graph window, with the graph of pressure vs. volume and its best-fit curve displayed.
Enter your name(s) and the number of copies you want to print.
Chemistry 125 Laboratory Manual –2004-2005
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Boyle’s Law
Data and Calculations
Volume(mL)
Pressure(atm)
Constant, k(P / V or P•V)
Questions:
1. If the volume is doubled from 5.0 mL to 10.0 mL, what does your data show happens to the
pressure? Show the pressure values in your answer.
2. If the volume is halved from 20.0 mL to 10.0 mL, what does your data show happens to the
pressure? Show the pressure values in your answer.
3. From your answers to the first two questions and the shape of the curve in the plot of pressure versus
volume, do you think the relationship between the pressure and volume of a confined gas is direct or
inverse? Explain your answer.
4. Based on your data, what would you expect the pressure to be if the volume of the syringe was
increased to 40.0 mL? Explain or show work to support your answer.
5. Based on your data, what would you expect the pressure to be if the volume of the syringe was
decreased to 2.5 mL? Explain or show work to support your answer.
6. What experimental factors are assumed to be constant in this experiment?
7. Using P, V, and k, write an equation representing Boyle’s law. Write a verbal statement that
correctly expresses Boyle’s law.
Chemistry 125 Laboratory Manual –2004-2005
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EXTENSION
1. To confirm the type of relationship that exists between pressure and volume, a graph of pressure
versus the reciprocal of volume (1/volume or volume-1) may also be plotted. To do this using
Logger Pro, it is necessary to create a new column of data, reciprocal of volume, based on your
original volume data.
• Remove the Linear Regression box from the graph by clicking on its upper-right corner.
• Choose New Column Formula from the Data menu.
• Enter “1/Volume” as the Long Name, “1/V” as the Short Name, and “1/mL” as the Unit. Then
click on the Definition tab.
• Enter the correct formula for the column (1/volume) into the Equation edit box. To do this, type in
“1” and “/”. Then select “Volume” from the Variables list. In the Equation edit box, you should
now see displayed: 1/“Volume”. Click OK .
• Click on the vertical-axis label, select “Pressure” (only), and click OK . Click on the horizontalaxis label, select “1/Volume” to be displayed on the horizontal axis, and click OK .
2. Decide if the new relationship is direct or inverse and change the formula in the Fit menu
accordingly.
• Click the Curve Fit button, .
• Choose Variable Power (y = Ax^n) from the list at the lower left. Enter the value of n in the
exponent edit box that represents the relationship shown in the graph (e.g., type “1” if direct, “-1”
if inverse). Click Try Fit .
• A best-fit curve will be displayed on the graph. If you made the correct choice, the curve should
match up well with the points. If the curve does not match up well, try a different exponent and
click Try Fit again. When the curve has a good fit with the data points, then click OK .
3. If the relationship between P and V is an inverse relationship, the plot of P vs. 1/V should be direct;
that is, the curve should be linear and pass through (or near) your data points. Examine your graph to
see if this is true for your data.
4. (Optional) Print a copy of the Graph window. Enter your name(s) and the number of copies of the
graph you want, then click OK .
Chemistry 125 Laboratory Manual –2004-2005
Page 98
Molar Volume of Hydrogen Gas
Goals:
Apply Dalton’s law of partial pressures to a mixture of gases.
Use the general gas law to calculate information about a gas.
Background:
When the volume of a gas is measured at a certain temperature and pressure, its volume can be
calculated at another temperature and pressure using the relationship below:
P1V1 P2V2
=
T1
T2
In this equation P is the pressure of the gas, V is the volume of the gas, and T is the absolute (Kelvin)
temperature. The subscript one refers to the initial state of the gas while the subscript two refers to the
final state.
In a mixture of gases, each gas exerts pressure independently of the presence of the other gases. The
total pressure of the mixture is the sum of the partial pressures of each gas:
Ptot = P1 + P2 + . . .
The result is known as Dalton’s law of partial pressures.
In this experiment you will be reacting an active metal (Mg) with H+ (acid) and collecting the hydrogen
gas (H2) which is produced. The gas will be collected over water so there will also be water vapor
present. Dalton’s law of partial pressures will be used to determine the pressure of the hydrogen gas
alone.
The moles of hydrogen gas should be equal to the number of moles of magnesium metal used, as can be
seen from the balanced reaction:
Mg + 2HCl → H2 + MgCl2
Table of Water Vapor Pressure at Various Temperatures: (from CRC Handbook)
17°
18°
19°
20°
21°
22°
14.5 mm Hg
15.5
16.5
17.5
18.6
19.8
23°
24°
25°
26°
27°
21.1
22.4
23.8
25.2
26.7
In this experiment:
Patm = PHydrogen gas + Pwater vapor + Pwater column
Chemistry 125 Laboratory Manual –2004-2005
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Procedure:
Obtain a small piece of magnesium, about 0.03 g (about 3 cm). Clean the strip of magnesium with steel
wool and wipe it clean. Weigh the magnesium and record the mass.
Curl up the magnesium strip and tie a string around it, leaving one end of the string about 5 cm long.
Obtain a eudiometer tube and a one-hole stopper that fits the tube. You will also need a 250 mL beaker
that is at least half filled with water.
Hold the eudiometer tube so that the closed end is at the bottom and clamp it to a ring stand. Pour about
10 mL of 6 M HCl into the tube. Slowly pour deionized water down the side - you do not want the acid
and water to mix yet! Fill the tube completely with water, past the markings to the top of the tube.
Put the string-wrapped magnesium at the top of the tube, leaving the long end of the string on the
outside of the tube. Insert the rubber stopper into the tube, catching the string against the stopper.
(Some water will be pushed out.) The tube should be full of water with the magnesium held close to the
stopper by the string. Place the 250 mL beaker (partially filled with water) under the tube.
Unclamp the eudiometer tube and turn it so that the stoppered end is down. Make sure this end is in the
250 mL beaker. Hold the tube in place with the clamp. Watch as the more dense acid falls down
through the water, reacting with the magnesium when the acid reaches the bottom of the tube.
Hydrogen gas will bubble up toward the top of the tube and water will run out of the hole in the stopper
into the beaker.
Allow the apparatus to cool when reaction is complete. Measure and record the volume of hydrogen gas
collected. Record the temperature of the room (since the apparatus has cooled, this will also be the
temperature of the gas). Measure the height in mm of the water remaining in the eudiometer tube.
The liquid in the beaker will be acidic. Neutralize the acid with baking soda and wash down the drain
with water.
Chemistry 125 Laboratory Manual –2004-2005
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Molar Volume of Hydrogen Gas
Data and Calculations
Name _________________________
Sec ___
Trial 1
Trial 2
Trial 3
Mass of magnesium metal
____________
____________
____________
Moles of Mg
____________
____________
____________
Moles of H2
____________
____________
____________
Volume of H2 gas
____________
____________
____________
Temperature (°C)
____________
____________
____________
Temperature (K)
____________
____________
____________
Height of water column (mm H2O)
____________
____________
____________
Height of water column (mm Hg)
____________
____________
____________
Total pressure (mm Hg)
____________
____________
____________
Partial pressure of water (mm Hg)
____________
____________
____________
Partial pressure of H2 (mm Hg)
____________
____________
____________
Volume of H2 at STP
____________
____________
____________
Molar volume of H2 at STP
____________
____________
____________
Calculation for the conversion of mm H2O to mm Hg:
Height (mmHg) = height (mm H2O) / 13.6
Show calculations for one trial.
Chemistry 125 Laboratory Manual –2004-2005
Page 101
Gas Law Questions
1. What type of chemical reaction occurs when magnesium react with hydrochloric acid?
Briefly explain your answer:
2. For the following errors, indicate whether the calculated value for the molar volume of H2 will be
too high, too low, or not affected. Explain your choice.
a. Some of the magnesium metal did not react with the acid.
Explain:
b. You forgot to subtract the vapor pressure of water from the total pressure of hydrogen gas.
Explain:
3. Standard temperature and pressure (STP) for gases is a reference set of conditions where the
temperature is 0ºC and the pressure is 1 atm.
a. Use the ideal gas law to calculate the molar volume of an ideal gas at STP to five significant
figures (R = 0.082057 L atm/mol K).
b. Calculate the percent error of your experimental molar volume of H2 at STP from the value for
an ideal gas.
4. All gases deviate to some extent from the behavior of perfect gases. At standard conditions, the
density of O2(g) is 0.0014290 g/mL, that of H2(g) is 0.00008988 g/mL, and that of CO2(g) is
0.0019769 g/mL. Using these values and the precise atomic weights, calculate the molar volume
of each of these gases, in L to five significant figures.
Chemistry 125 Laboratory Manual –2004-2005
Page 102
Appendix A: Statistical Treatment of Data
Accuracy:
One measure of the accuracy of the result is the percent error. To calculate the percent error, one
compares the experimentally determined value of some quantity to its accepted value. The following
formula is used:
% error =
(experimental - accepted)
x 100
accepted
Example: Suppose an experimental determination of the density of copper gives the result, d= 8.77
g/cm3. The accepted value is 8.95 g/cm3. Calculate the percent error.
% error =
(8.77 - 8.95)
8.95
x 100 = - 2.0 %
The negative sign is the result above indicates that the experimental value is 2.1% lower than
the accepted value.
Precision:
When an experiment is repeated several times so that the same quantity is determined several times, it is
common practice to calculate a statistic that indicates how well the individual trials agree with one
another. For larger data sets, the standard deviation is often calculated assuming that the error in the
individual results is randomly distributed. If the data is precise, the standard deviation will be relatively
small.
For smaller data sets (3-5 trials) it is better to calculate the relative average deviation (r.a.d.) of your
data. The r.a.d. is a measure of the precision of your data. Like the standard deviation, it does not
indicate the accuracy of your result, but it does provide an indication of the reproducibility of the results.
A set of data, which is precise, has a relatively low r.a.d.
The r.a.d. in parts per thousand is defines as follows:
r.a.d. =
average deviation
mean
x 100
_
∑ xi
where, x = mean =
n
_
∑| xi - x |
and, average deviation =
n
where n= the number of data points and xi= the individual data points
Note: the summations in the formulae above go from i = 1 to n.
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Example:
For the determination of the percent of iron in an ore, the following data were obtained:
22.24%, 22 29%, 22.61%, and 22.52%. Calculate the relative average deviation of the data.
Mean = (22.45 + 22.29 + 22.61 + 22.52)/4 = 22.4675
The average should be reported as 22.47%. Extra digits (indicated above by the underlining) are
temporarily carried here to avoid round off error in subsequent calculations).
Deviations:
[22.45 – 22.4675] = 0.0175
[22.29 – 22.4675] = 0.1175
[22.61 – 22.4675] = 0.1425
[22.52 – 22.4675] = 0.0525
Average deviation = (0.0175 + 0.1175 + 0.1425 + 0.0525)/4 = 0.0825
r.a.d.= (0.0825/22.4675) x 100 = 0.37 %
The r.a.d. for this example indicates a precision of 0.37 parts out of one hundred for the data.
When calculating the relative average deviation, be sure to carry out some "extra digits" through the
calculation to avoid round off error. Round off your final answer to the appropriate number of
significant digits.
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Equipment Hunt
Name ______________________________________
Sec _____
Write the name of each piece of equipment beside its picture. You may need to look for equipment in
places other than your equipment drawers.
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II. Make a rough sketch of each of these pieces of equipment.
Tweezers
Iron ring
Ring stand
Test tube rack
Beaker tongs
Spatula
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III. Some application questions.
1. Which piece of equipment would be best for
- Carrying a hot beaker from one place to another.
- Carrying out a reaction with small amounts of reagents and observing the results
- Evaporating liquid solvent
- Measuring 50.0 ml of liquid as precisely as possible
2. What would be a good use for
- a spatula
- a clamp
- a test tube rack
- a funnel
3. What sizes are the graduated cylinders in the equipment drawer that you use?
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