2015 - 2016 Version
Table of Contents
Principles for Safety in the Chemical Laboratory. . . . . . . . . . . . . . . . . . . 1
Laboratory Notebooks. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 6
Experiment 1. Chemical Measurements and Significant Figures. . . . . . 8
Experiment 2. Percent Composition of Metal Oxides .. . . . . . . . . . . . . . 13
Experiment 3. Atomic Spectroscopy. . . . . . . . . . . . . . . . . . . . . . . . . . . . . 18
Experiment 4. Spectrophotometry.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 24
Experiment 5. Identification of Unknown Solutions. . . . . . . . . . . . . . . . 38
Experiment 6. Molecular Models.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 42
Experiment 7. Isomers, Hybridization, and Molecular Orbitals. . . . . . 49
Experiment 8. Copper Compounds. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 62
Experiment 9. Synthesis of a Cobalt Salt. . . . . . . . . . . . . . . . . . . . . . . . . 70
Experiment 10. Reaction Stoichiometry. . . . . . . . . . . . . . . . . . . . . . . . . . 75
Experiment 11. Redox Titration. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 82
Experiment 12. The Ideal Gas Law. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 90
i
Experiment 13. Molar Mass of a Vapor. . . . . . . . . . . . . . . . . . . . . . . . . . 96
Experiment 14. Thermochemistry. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 102
Experiment 15. Determination of Glucose using a Spectrophotometer
. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 111
Experiment 16. Physical Properties of Chemicals - Melting Points,
Boiling Points and Sublimation. . . . . . . . . . . . . . . . . . . . . . . . . . . 116
Experiment 17. Determination of the Enthalpy of Vaporization of H2O
. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 134
Experiment 18. Freezing-Point Depression. . . . . . . . . . . . . . . . . . . . . . 139
Experiment 19. Rate Law for the Iodine Clock Reaction. . . . . . . . . . . 145
Experiment 20. Reaction Rates. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 152
Experiment 21. Solubility of Calcium Iodate . . . . . . . . . . . . . . . . . . . . 158
Experiment 22. Acid-Base Strength of Salts.. . . . . . . . . . . . . . . . . . . . . 165
Experiment 23. Buffers and Potentiometric Titrations . . . . . . . . . . . . 174
Experiment 24. Structures of Organic Molecules. . . . . . . . . . . . . . . . . 178
Appendix 1. Data Analysis. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 183
Using Excel to Graph Data (Office 2013 version). . . . . . . . . . . . . . . . . 191
Using Excel to Graph Data (Office 2007 version). . . . . . . . . . . . . . . . . 197
Constants. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 207
ii
PREFACE
You cannot truly appreciate any science without getting your hands wet by doing
experiments. When I first came to Black Hills State in 1998, I found a record of several
experiments that had been done in the past, but no organized lab manual that put all the
experiments together into a single cohesive book. I had two choices, either I could
continue to come up with single experiments on an ad hoc basis and give them to the
students week by week, or I could write a lab manual that contained many useful
experiments in a single organized volume.
Since I knew the task of writing a lab manual could consume days and weeks (months,
years) of my time, I did what any sane person would do, I looked at somebody else’s
manual for ideas and inspiration. In particular, since I had just taught at Ohio Northern
University and knew that they had a manual full of experiments that worked well in a 2-3
hour lab period, I asked their permission to copy their manual and use it here at Black
Hills State.
Fortunately they allowed me to do this, and, in 1998 at least, most of the experiments in
this manual are almost direct copies taken from the Ohio Northern University
Introductory Chemistry Lab Manual. As such, I must acknowledge the work and effort of
the staff from that institution, and I thank them whole-heartedly for letting me use their
work here.
The story does not stop there. Making a lab manual is an evolutionary process. Things
that worked at one university will work differently at another because there is a different
physical set-up to the lab and a different way of doing things. Further, some experiments
I like as they are, while others I want to change in some way to make them work better. I
will also continually be looking for other experiments that might illustrate a given idea
better. Thus, as each year passes the manual will change and evolve. At this beginning of
this process I want to acknowledge the help of Jennifer Zoller who has spent a great deal
of time helping me to put this manual together.
iii
Principles for Safety in the
Chemical Laboratory
Safe practices in the chemical laboratory are of prime importance. A student should
consider it an essential part of his or her educational experience to develop safe and
efficient methods of operation in a lab. To do this, one must acquire a basic knowledge of
properties of materials present in the lab, and one should realize the types of hazards that
exist and the accidents and injuries that can result from ignorance or irresponsibility on
the part of the student or a neighbor.
Regulations
1. Wear safety goggles at all times while in the laboratory.
2. Report all accidents to the instructor or lab assistant immediately.
3. NEVER eat, drink, chew, or smoke in the laboratory.
4. NEVER leave an experiment unattended. Inform the lab assistant if you must leave
the lab.
5. After the experiment is completed, turn all equipment off, making sure it is properly
stored, and clean your area.
Failure to comply with these regulations is cause for immediate dismissal from lab.
Precautions
1. Approach the laboratory with a serious awareness of personal responsibility and
consideration for others in the lab.
2. Become familiar with the location of safety equipment, such as acid-base neutralizing
agents, eye wash, fire extinguisher, emergency shower, and fire blanket.
3. Pay strict attention to all instructions presented by the instructor. If something is not
clear, do not hesitate to ask the instructor or lab assistant.
4. Clean up all chemical spills immediately.
1
5. Be aware of all activities occurring within a reasonable proximity of yourself since
you are always subject to the actions of others.
6. To avoid contamination of community supplies, do not use personal equipment such as
spatulas in shared chemicals and replace all lids after use.
7. Avoid unnecessary physical contact with chemicals; their toxic properties may result
in skin irritation.
8. Use all electrical and heating equipment carefully to prevent shocks and burns.
9. NEVER handle broken glassware with your hands; use a broom and a dust pan.
10. Wash your hands at the end of the laboratory.
Personal Attire
Choice of clothing for the laboratory is mainly left to the discretion of the student.
Because of the corrosive nature of chemicals, it is in your best interest to wear
comfortable, practical clothing. Long, floppy sleeves can easily come into contact with
chemicals. A lab coat is suggested to help keep clothes protected and close to the body.
Accessories also need consideration. Jewelry can be ruined by contact with chemicals.
Open toed shoes do not adequately protect one against chemical spills. If hair is long
enough to interfere with motion or observation, it should be tied back. Remember that
your clothes are worn to protect you.
Assembling Equipment
Equipment should be assembled in the most secure and convenient manner. Utility
clamps are provided to fasten flasks, etc., to the metal grid work located at the center of
each bench. This keeps top-heavy or bulky equipment away from the edge where it can be
knocked easily off the bench. Consider the safe location of the hot plate. Keep it near the
grid work to minimize chances of contact with the body. If the aspirator is being used,
locate your apparatus near the sink for convenience.
2
Handling Glassware
Laboratory glassware is usually fragile, and if it is not properly handled, serious injuries
may result Do not force glass tubing or thermometers into a rubber stopper. Lubricate the
tubing or thermometer with glycerol or water, wrap it in a towel, and gently insert it into
the stopper by using pressure in a lengthwise direction while rotating it. Always grasp the
tubing near the stopper. When removing the tubing, remember to protect your hands with
a towel. If there are difficulties with this procedure, ask for the instructor's assistance.
Apparatus that can roll should be placed between two immobile objects away from the
edge of the bench. Chipped or broken glassware cannot be used. There are special
receptacles near each bench for these waste materials. After the experiment is completed,
all glassware should be emptied, rinsed, and cleaned.
Acids and Bases
In this lab sequence, you will come in contact with several acids and bases. As with all
chemicals, caution must be taken to prevent contact with the skin. When handling these
chemicals, keep hands away from the eyes and face until they have been thoroughly
washed. If an acid or base comes in contact with your skin, flush the area with large
quantities of clean, cold water. Eyes are extremely sensitive. Use the eye wash provided
in the laboratory, or wash with water for at least 10 minutes. Again, the instructor must be
notified immediately. To insure your safety, neutralize acid or base spills before cleaning
them up. Boric acid solution is available to neutralize base spills, and carbonate powder is
provided to neutralize acids.
3
Attention:
Students are advised against wearing contact lenses while
observing or participating in science laboratory activities. While
hard contact lenses do not seem to aggravate chemical splash
injuries, soft contact lenses absorb vapors and may aggravate
some chemical exposures, particularly if worn for extended
periods.
Please take your contact lenses out prior to entering the
laboratory.
Contact Lens Administrative Policy and Waiver Form
Students are advised against wearing contact lenses while observing or
participating in science laboratory activities. While hard contact lenses do not
seem to aggravate chemical splash injuries, soft contact lenses absorb vapors
and may aggravate some chemical exposures, particularly if worn for
extended periods. You are asked to please remove your contact lenses prior
to entering the laboratory.
If you do not wish to comply with this recommendation, you must fill out the
next page, which is a waiver form.
4
Waiver of Liability, Indemnification and Medical Release
I am aware of the dangers involved in wearing contact lenses in a science
laboratory setting. On behalf of myself, my executors, administrators, heirs,
next of kin, successors, and assigns, I hereby:
a. waive, release and discharge from any and all liability for my personal
injury, property damage, or actions of any kind, which may hereafter, accrue
to me and my estate, the State of South Dakota, and its officers, agents and
employees; and
b. indemnify and hold harmless the State of South Dakota, and its officers,
agents and employees from and against any and all liabilities and claims
made by other individuals or entities as a result of any of my actions during
this laboratory.
I hereby consent to receive any medical treatment, which may be deemed
advisable in the event of injury during this laboratory.
This release and waiver shall be construed broadly to provide a release and
waiver to the maximum extent permissible under applicable law.
I, the undersigned participant, acknowledge that I have read and understand
the above Release.
Name _______________________________ Age _________________
Signature ____________________________ Date _________________
Section _____________________
Is there any health information you would like us to know if there is an
accident?
5
Laboratory Notebooks
You are required to use a bound notebook in CHEM 112-114 lab to record all primary
data and observations. You should prepare your notebook each week before coming to lab
by writing the title of the experiment on a new numbered page, summarizing relevant
equations from the lab manual, and starting calculations involving molar masses, etc.
Take note of theoretical ideas and special instructions given by your instructor at the start
of each experiment. Your notebook should be a complete record of your work in lab. You
or other chemists should be able to understand the notes in the future, not just during the
current experiment Good note taking in lab is a valuable skill that you can learn with a
little effort and practice.
Guidelines to be Followed:
1. Always bring your notebook with you to lab. You will be graded on the completeness
of your previous note taking and your preparation for the current experiment. You may
use your notebook during a lab quiz.
2. Number the pages sequentially and reserve space at the beginning for a table of
contents.
3. Take your notebook to the balance room, etc. and record values directly in it - not on
loose scraps of paper.
4. Specify each measured quantity by name and include the units.
5. If you make a mistake in your notebook, simply draw a solid line through the error and
write the correction nearby.
6. Tables greatly simplify data entry; they should be set up before coming to lab.
7. Write down all observations such as color and phase changes - don't rely on your
memory.
8. Save time by doing trial calculations in your notebook before filling out any report
sheets.
9. Save time by making preliminary sketches of graphs on the ruled lines in your
notebook.
10. Make observations of what you actually see, not what you think you should see!
6
10. Diagrams of experimental set-ups. You will be putting together equipment that you
have never seen before; you need some way to remember how you put the apparatus
together.
11. Conclusions. First think about the purpose of the lab. What is it that you were trying
to accomplish? Now write up a paragraph summarizing how you accomplished that
purpose. What were the key experiments, observations or calculations that allowed you
to accomplish the stated purpose of the lab?
7
Experiment 1.
Chemical Measurements and Significant Figures
Purpose: In this experiment you will be:
i
taking various measurements using equipment in the lab
i
evaluating significant figures
i
doing some calculations with significant figures
Note: The individual parts of the lab do NOT have to be done in a particular order, so
feel free to do any measurement or calculation at any time, as the equipment is free for
use.
I. Measuring mass - Use of a the balances.
Each person will get an object of unknown weight. Your job will be to determine the
weight of that object on each balance we have in the lab. Please remember that I have
already weighed the object, and that if you touch it with your fingers you will leave oils
behind that will change the weight of the object. Thus you should never handle the object
with your hands. You may use tongs or a piece of paper, but never your hands.
Also, you should treat this object like a chemical, that is, do not simply put it on that
balance pan. You should first put a piece of paper on the balance and either record its
weight or “tare” the weight on the balance. Then you place your object on the paper and
get the total weight of object and paper. The actual weight of the object will be the
difference between the weight of the (object + paper) minus the weight of the paper.
As you use each balance you will see that they have a level of uncertainty. Some can
measure to 1g, others to .0001g, and others in between. Pay attention to the inherent
accuracy of each balance, for when you report your numbers on the write up, and for
future reference, when you need to decide on your own how to measure the amount of a
particular chemical.
II. Measuring volumes
The chemist has many different ways to measure volumes. Lets look at some of them and
think about the level of uncertainty in the measurement.
A. Beakers - Probably the crudest method is to simply pour the water into a
beaker. Find beakers A, B, and C and record the volume of water in those beakers on the
reporting sheet using the appropriate number of significant figures.
8
B. Graduated cylinders - The next best way to measure a volume is with a
graduated cylinder. This long skinny tube has many more calibration marks than a
beaker, and this makes it easier to determine volume with a higher precision. Find
Graduated cylinders A, B, C, and D. and record the volume of water in these cylinders
with the appropriate number of digits on the reporting sheet
C. Volumetric Flasks - If you want a solution to contain a specific total volume,
you usually make the solution in a volumetric flask. This flask has been calibrated to
contain a precise amount of liquid when filled to a line on the neck of the flask. The
hardest part of using a volumetric flask is getting the liquid level to hit the line exactly.
Take a volumetric flask and fill it to the line. When this is done have the lab instructor
initial the reporting sheet. Note that since you fill a volumetric flask with water, it does
not have to be dry when you start filling it. Thus don’t worry about trying to get the flask
dry before you start.
Some flasks are marked with the uncertainty in the volume contained, while others have
no marking. If there is no marking, the uncertainty is considered to by ± 1 drop or about
0.05 ml
D. Volumetric pipets - These are pipets with large bulbs that have been calibrated
with a mark or line, so when they are filled to the mark, they contain a given volume.
Like the volumetric flask the biggest problem is getting the liquid exactly to the line.
Please remember that you never use your mouth for pipeting. Using one of the pipet
bulbs provided, fill a volumetric pipet to the calibration line and have the instructor initial
your reporting sheet. Like the volumetric flask the uncertainty in a volume delivered by a
pipet is assumed to be ± 1 drop (0.05 ml) unless otherwise marked on the pipet.
E. Burets - While volumetric flasks and pipets can deliver very accurate volumes,
they have no flexibility and can only deliver one fixed volume. How can you deliver
volumes in a manner that is both flexible and accurate? The answer to this is the buret. A
long tube with fine calibration marks and a stopcock that allows you to deliver and
particular volume you wish. The problem most people have with a buret is reading it
accurately. Three burets have been filled by the instructor. Find the burets, read the
liquid levels in the burets, and write them down on the reporting sheet with the
appropriate level of precision. Remember that you should be able to interpolate the
reading to one more decimal point than the buret is calibrated.
III. Significant Figures
Review your text on significant figures and propagation of error in calculations and
answer the problems on the report sheet.
9
Name:
Report Sheet
Chemical Measurements and Significant Figures
I. Measuring mass
Object Number______________
Object weight (remember to report the correct number of significant digits)
Electronic Balances
Balance A
Weight of Object (g)
________±______
Balance B
________±______
Balance C
________±______
Balance D
________±______
Pan Balance
Object + Paper
Balance E
Paper
Object
__________±____ ___________±_____
________±______
II. Measuring volume
A. Beakers
Record the volumes of water in the beaker with the appropriate number of
significant figures.
Raw Volume (estimated) ± uncertainty
Volume
(Using sig figs)
Beaker A
________________
____________
_________
Beaker B
________________
____________
_________
Beaker C
________________
____________
_________
10
B. Graduated Cylinders
Record the volumes of water in the graduated cylinders
Raw Volume (estimated)
± uncertainty
Volume
(Using sig figs)
Cylinder A
________________
____________
__________
Cylinder B
________________
____________
__________
Cylinder C
________________
____________
__________
Cylinder D
________________
____________
__________
C. Volumetric Flasks
Flask filled _____________ (instructors initials)
Volume and uncertainty in volume? _______________________
D. Volumetric Pipets
Pipet filled _______________(instructors initials)
Volume and uncertainty in volume? _______________________
E. Burets
Liquid level observed (raw)
Uncertainty in level
A.
_____________
______________
Reading
(appropriate sig.
fig.)
_______________
B.
_____________
______________
_______________
C.
_____________
______________
_______________
Question - If A. was your initial buret reading and B was the reading after you delivered
some of the liquid into a flask, how much liquid did you put in that flask?
Question - In the answer above, what is the uncertainty in the delivered volume?
11
III. Significant Figures
1. Indicate the number of significant figures in the following measurements.
(a) 0.209 mL _____________
(b) 0.00077 g _______________
(c) 135.7 g
(d) 21.5 mL _______________
_____________
(e) 0.0302 g _____________
(f) 1.020 g/mL _______________
2. The following represent results of mass and volume measurements used to determine
the density of some liquids. Calculate each density to the appropriate number of
significant figures.
(a) 12.5 g / 9.5 mL
________________
(b) 1.049 g / 10.00 mL
________________
(c) (22.892 g - 4.3380 g) / (23.45 mL - 0.05 mL)____________________
IV. Calculator Worksheet
Perform the following calculations on your calculator. If your answer does not match the
answer given, consult your laboratory instructor.
A.
2+3x2=8
B.
4 x 10 – 2 / 5 = 39.6
C.
5 x 10-2 = 0.05
D.
5 x 10-2 – 2 = -1.95
E.
(5 x 10-2)2 = 2.5 x 10-3
F.
2 x 3 x103 + 2 / 8 = 6000 or 6 x 103
G.
3 x (9 x 10-3)5 = 1.77x10-10
12
Experiment 2.
Percent Composition of Metal Oxides
Purpose:
i
i
i
In this experiment you will use the reaction, Mg(s) + O2(g) 6MgxOy (s) to
determine the % magnesium and % oxygen in magnesium oxide.
You will then observe a similar but opposite reaction, AgxOy(s) 6 Ag(s) +
O2(g) to determine the % silver and % oxygen in silver oxide.
Then, by comparing the experimental % compositions to the theoretical %
composition of known formula, you will determine the molecular formula
of the two oxides
Background
Many metals react with oxygen to from oxides. Some metals, like Magnesium, do this
extremely well, and can burn with just the O2 in the air and a little heat to get the reaction
started. Other metals, like iron are less reactive, and won’t burn easily, but do slowly
oxidize when exposed to air and moisture (Iron rusting). Still other metals, like gold,
form oxides only under the most extreme conditions.
In this lab you will look at two metal oxides, silver oxide and magnesium oxide.
Magnesium is a highly reactive metal, and burns vigorously once you get it hot enough to
ignite. In fact, it is so reactive, that a second side reaction begins to occur in which the
magnesium reacts with the relatively inert N2 gas in the air: 3Mg(s) + N2(g) 6Mg3N2(s).
In this experiment we will burn the magnesium in a covered crucible to control the rate of
the reaction, and then do a second reaction, Mg3N2(s) + 6H2O(l) 6 3Mg(OH)2(s) +
2NH3(g) to convert any of the magnesium nitride to magnesium hydroxide. The
magnesium hydroxide formed in this reaction is then converted into magnesium oxide by
simple additional heating.
From the mass of magnesium you start with, and the mass of magnesium oxide that
remains after the reactions, you can calculate the % composition of magnesium and
oxygen in magnesium oxide.
Silver is a metal that does not oxidize easily at room temperature and, in fact, the oxide of
silver decomposes back into silver and oxygen if it is simply heated. In this part of the
lab your instructor will start with silver oxide, heat it to the temperature that it
decomposes, and isolate the silver that remains after this reaction. By comparing the
mass of the silver oxide you started with and the mass of silver that remains after the
sample is heated you can calculate how much oxygen was in the silver oxide, and from
this you can calculate the % composition of silver and oxygen in silver oxide.
13
Procedure:
Formation of Magnesium Oxide
To prove that magnesium readily forms oxides the instructor will ignite a 0.1 gram piece
of magnesium ribbon. Record your observations in your lab notebook. (Note: The
magnesium flare is so bright, you should not look directly at it when it ignites.)
Now that you have seen the uncontrolled reaction, you can try the reaction under a more
controlled environment.
Wash and dry a crucible and lid. Remember once you have washed the crucible & lid
use your crucible tongs to handle it NOT YOUR FINGERS!! Place a ceramic triangle
on a ring stand and position your propane torch about 3 inches under the triangle. Place
your cleaned crucible and lid on the triangle, light your propane torch, and heat your
crucible and lid for about a minute to evaporate any remaining moisture. Turn off your
burner and let the crucible and lid cool for about 10 minutes.
After the crucible and lid are cool take them to the balance and weigh them to the nearest
.001 g. and record this data in your lab notebook. Next, obtain a piece of magnesium
ribbon about 1.5 inches long. Cut this ribbon into 4-5 pieces and place a kink in the
middle of each piece. Place these ribbon pieces in the crucible, and accurately weigh the
crucible, lid, and the ribbon to the nearest .001g and record this data in your notebook.
Also record the appearance of magnesium.
Place crucible back on the triangle and cover with lid slightly ajar to allow air to circulate.
Turn on the torch and heat the crucible and lid for 10 minutes over the hottest part of the
flame. Adjust the flame under the crucible so that the bottom of the cricible glows, but
not the sides. At the end of the ten minutes turn off the burner and allow the crucible to
cool until you can almost touch it. Use your crucible tongs to remove the lid and examine
the contents of the crucible. What do you see? Record your observations in your lab
notebook.
While the crucible is cooling, fined the deionized water that is heated to boiling on the
hotplate. Add 20 drops of this boiling hot water to the warm crucible. Note any odor
that is released when you add the water.
Now put the crucible lid back on (again slightly ajar) and gently heat the crucible for 10
more minutes to evaporate the water and complete the conversion of Mg2N3 to
magnesium oxide.
14
If the reaction looks complete, use your crucible tongs to put the crucible in your
casserole dish, cover with the lid, and let cool for 10 minutes. Once cool, weigh the
crucible, lid and contents and record the final weight in your lab notebook.
Clean, dry and re-weigh, the crucible and lid, before repeating the experiment. You will
need data from a total of two runs of the experiment.
Decomposition of Silver oxide
The instructor will accurately weigh a clean, dry, crucible & lid. Record this value in
your lab notebook. Then the instructor will accurately weigh about 0.5 g of silver oxide
into the crucible. (Again record this weight, and the appearance of the silver oxide in
your notebook.) The instructor will then heat the crucible, silver oxide, and lid with a
propane torch. To observe the reaction, the instructor may leave the crucible lid off. At
that point the lid should be placed on the crucible so all the ‘soot’ is retained for a proper
weight. When the reaction is complete (~ 10-15 minutes) the instructor will carefully
remove the lid to see if the reaction looks complete. If it is, then the crucible and lid will
be moved to a casserole dish to cool. Once cooled the crucible, lid and product will be
weighed. The instructor will write the mass of crucible, lid and product on the board.
Record this value in your lab notebook so you will have all the raw data necessary to
determine % Ag and %O in the original oxide.
Calculations:
Using the data from your three runs of the reaction of magnesium with oxygen,
determine the mass of magnesium and the mass of magnesium oxide. Then calculate the
mass of oxygen and the %Mg and %O in the magnesium oxide. By comparing the
experimental %Mg and %O to the theoretical % composition of known formulas,
determine the molecular formula for the magnesium oxide formed. Write a balance
chemical equation for the formation of magnesium oxide from magnesium and oxygen.
For the silver oxide experiment use the instructor’s raw data to determine the mass of
silver oxide used, and the mass of silver produced. Then calculate the mass of oxygen in
the oxide produced. Calculate the %Ag and %O. If the instructor has data from multiple
runs, determine the average %Ag and %O from these trials and compare to the table of
theoretical % compositions of known formulas to determine the molecular formula of the
silver oxide. Do these calculations in your lab notebook and transfer the information to
the Report Sheet to turn in. Also write a balanced chemical equation for the reaction for
the decomposition of silver oxide into silver and oxygen.
Make sure that you have all of the calculations set up and worked in your lab notebook.
Also make diagrams of the apparatus used in these two experiments in you lab notebook.
15
Name: ___________
___________
Report Sheet
Determination of the Molecular Formula of Silver Oxide and Magnesium
Oxide from Experimental % Composition
I. Oxidation of Magnesium
Mass of crucible, lid, and magnesium
Trial 1
________
Trial 2
________
Mass of empty crucible and lid
________
________
Mass of magnesium used in reaction
________
________
Mass of magnesium oxide, crucible & lid
________
________
Mass of empty crucible & lid used
________
________
Mass of magnesium oxide formed:
________
________
Mass of oxygen in magnesium oxide:
________
________
% O in magnesium oxide:
________
________
Average ______
% Mg in magnesium oxide:
________
________
Average ______
Calculate the theoretical % magnesium and % oxygen for the following molecular
formulas
MgO
Mg2O
MgO2
% Mg
%O
What is your best guess for the formula of Magnesium oxide?
Why?
Write a balanced chemical equation for the formation of magnesium oxide from
magnesium and oxygen:
16
II. Decomposition of Silver Oxide
Trial 1
________
________
________
________
________
________
Trial 2
(If available)
________
________
________
________
________
________
Trial 3
(If available)
________
________
________
________
________
________
mass of oxygen in silver oxide
_______
_______
________
% Silver in silver oxide
_______
_______
_______
Mass of crucible, lid & silver oxide
Mass of empty crucible & lid
Mass of silver oxide
mass of crucible, lid & silver
mass of empty crucible & lid
mass of silver
_______ Average % Silver
% Oxygen in silver oxide
_______
_______
_______
_______ Average % Oxygen
Calculate the theoretical % silver and % oxygen for the following molecular formulas
AgO
Ag2O
AgO2
% Ag
%O
What is your best guess for the formula of Silver oxide?
Why?
Write a balanced chemical equation for the decomposition of silver oxide into
silver and oxygen:
17
Experiment 3.
Atomic Spectroscopy
Purpose: In this experiment you will:
i
analyze the spectrum of atomic hydrogen to determine
the Rydberg constant
i
use flame tests to identify unknown solutions
i
identify an unknown gas by its spectral "fingerprint"
Background
Refer to chapter 7 in your text and the next chapter, “Expt.4: Spectrophotometry" for
more background information.
Atoms in the gas phase have quantum states with very specific energies. The atoms each
exist in one of these states and may change to another state by absorbing or emitting a
quantum of light (i.e. photon). These "transitions" between quantum states occur at
distinct wavelengths (ë). Since atoms of each chemical element have a unique set of
quantum states, the observed wavelengths of the transitions constitute a spectral
"fingerprint" for that element While accurate measurements of the wavelengths are
usually the best way to identify the spectrum of an element, the intensity (i.e. brightness)
of the transitions is also a unique, qualitative feature that often aids in identification.
Energy sources such as electric discharges or flames excite atoms into high-energy
quantum states, and light emission may then be observed. When an aqueous salt solution
is vaporized in-a flame, metal cations are reduced to neutral atoms and the atoms are then
excited. Some metallic elements thus emit a characteristic visible color that may be used
to distinguish them in a set of unknown samples. The metals and corresponding flame
colors considered in this experiment are:
sodium
bright yellow
potassium
faint violet
strontium
red - orange
barium
yellowish green
lithium
deep red
copper
green/blue
18
Procedure
Experiment 1:
Determining the Rydberg Constant
Work in pairs for this part of the experiment. Your instructor will demonstrate the use of
a spectroscope to observe the emission spectrum of atomic hydrogen. Use the
spectroscope to measure the wavelengths in nm (3 s.f.) of as many spectral “lines” as you
can see. Refer to the provided emission spectrum of atomic hydrogen to confirm your
observations. Make a table of the wavelengths (ë) and convert them to wavenumbers (í)
using the relation
This set of lines is called the Balmer series. The lines correspond to transitions from
different upper quantum states with quantum numbers nu = 3, 4, 5, 6, etc. to a common
lower state whose quantum number is n1 = 2. The wavenumbers of the transitions are
given by the Rydberg formula,
where R is called the Rydberg constant.
Use figure 4.26, page 125 in your text to identify nu for each of the lines in your table. For
each of your observed line assignments (nu and í), use equation (2) to calculate a value
for the Rydberg constant, R, in wavenumber units (cm-1). Average the results for R, and
answer the related questions on the report sheet.
Experiment 2:
Flame Tests
Obtain a small sample (~1 mL) of each of the six unknown solutions in labeled test tubes.
Set up a propane torch and adjust it to give a stable, blue flame with an inner cone. Do
not leave the flame unattended. Use a clean wire loop to test each solution separately in
the flame. Carefully rinse the loop with distilled water between samples. Note the
distinctive color emitted by each sample as it is being vaporized, and match it to one of
the metals listed in the background information. Identify the metal in each unknown on
the report sheet
19
Experiment 3:
Emission Spectra of Gases
A spectral lamp containing an unknown gas will be set up for viewing through a simple
spectrometer. Notice the pattern of colored "lines” that correspond to the emission
wavelengths that have been dispersed. Sketch the pattern on the sheets provided and place
them in your notebook, including the color of each line.
Refer to the attached plots of visible emission spectra for several gases. Match the pattern
of colored lines that you observed with one of the provided spectra. Identify the unknown
gas on the report sheet.
Emission Spectra of Some Gases
Violet
Blue
Green
Yellow
20
Orange
Red
21
Name:
Report Sheet
Atomic Spectroscopy
I. Rydberg Constant (R)
Lower quantum number n1
Measured
Wavelength (nm)
Wavenumber
(m -1)
Upper quantum
number: nu
R
(m -1)
1. Show how you calculated R for you first line:
2. Your average value for R
3. Your standard deviation of R values
4. Does your average value for R agree with the accepted value, 1.097 x 107 m-1,
within ± 1 standard deviation? Suggest method(s) that could be used to improve the
agreement.
22
II. Flame Tests
1. Identify each of the unknowns by completing the following table:
Compounds
flame color
unknown #
SrCl2
BaCl2
NaCl
KI
CuCl2
LiOH
2. Estimate (not calculate) the emission wavelengths of Li, Na, and Ba atoms by matching
the flame colors that you observed with the spectral scale shown below. Which of these
types of atoms appears to emit photons with the lowest energy?
3. Write the electron configuration for the ground state of K atoms:
K atoms in an excited electron configuration, [Ar] 5p1, cause the violet
emission that you observed. K atoms may also be excited to the [Ar] 4p1
configuration. Is the 4p164s1 transition at a longer, shorter, or the same
wavelength as the 5p164s1 transition?
III. Emission Lines
Sketch the spectrum that you observed.
Unknown gas #_____ Identity___________
23
Experiment 4.
Spectrophotometry
Purpose: This lab is designed to introduce you to the basics of measuring how light
interacts with matter. Topics to be covered include:
•
Why materials appear to be different colors.
•
How the transmittance and adsorption of light varies with the amount of material
in the path of the light.
•
The origin of Beer’s Law.
•
How amounts of chemicals can be quantified using Beer’s Law.
Background
This lab is designed to help you discover the concepts behind the lab as you
perform the experiments. Thus I will not give you a background, you must discover these
things on your own! What you do need to know is how to run the spectrophotometer, the
instrument that you use to acquire the data for this lab.
Experimental Procedure
1.
If the spectrometer has a power cord, plug this cord into the socket. (This is
only for Red Tide USB 650 Spectrometers.)
2.
Connect SpectroVis Plus or Red Tide spectrometer to the computer using
the USB cable provided. If you have a SpectroVis Plus spectrometer (No
power cord), you can go on to step 3. If you have a Red Tide spectrometer
(Power cord) it may take a minute or two for the computer to recognize the
spectrometer and install the proper USB device. Wait for it.
3.
Now Start Logger Pro. Once Logger Pro starts, you should see a screen
with a bright rainbow.
a.
Look at the X axis. This tells you the wavelength of light, in
nanometers, of each of the colors you observe.
b.
Go to the report sheet and report the wavelength of the colors blue
green yellow and red.
c.
Look at the Y axis. If the Y axis is not in % transmittance take the
following steps:
Click on ‘Experiment’-‘Change Unit’-‘Spectrometer 1’ - ‘%
Transmittance’ The Y axis should now read Transmittance
(%).
24
What is % transmittance? % transmittance is a measure of how much light is
reaching the detector relative to how much light reaches the detector when there is
nothing in the light beam to absorb the radiation. 100% transmittance means that 100%
of the light reaches the detector (or 100% of the light is transmitted through the material).
0% transmittance means that 0% of the light reaches the detector (or 0% of the light is
transmitted through the material).
Before we can measure % transmittance, we need to calibrate the spectrometer so
the computer knows what the levels on its sensors correspond to 100% transmittance (All
of the light hitting the sensor) and 0% transmittance (no light hitting the sensor).
We will do this now. Check that there is nothing in the sample holder - the square
hole in the middle of the spectrometer box. Once you have emptied the sample holder
take the following steps:
•
Click ’Experiment’ - ‘Calibrate’ - ‘Spectrometer 1’
•
You should see the light turn on, and a window comes up saying that
the spectrometer must warm up for 90 seconds before it is properly
calibrated. Don’t skip this step. When the warmup is compete, the
window will say ‘Put a blank Cuvette in the device’.
•
Since we will not be using a cuvette in our first experiment, do not
put anything in the device and click on the ‘Finish Calibration’
button.
•
Click on the ‘OK’ button. The machine is now calibrated.
Experiment 1. Transmittance of light and color
At your lab station you should have three different colors of plastic sheets. Now,
remembering that 100% transmittance means that 100% of the light is transmitted through
a material, and by knowing the wavelengths that the different colors correspond to,
predict the transmittance spectra of your three different sheets of plastic on your
report sheet and have your instructor initial your predictions before you go any
further.
Once you have made your predictions, do the experiment.
•
Cut a piece of plastic from the sheet about 1 cm wide and place it in
the sample holder of the spectrometer in front of the light source.
•
Click on the Green ‘Collect’ button at the top of the screen.
•
You should now have a red line across your spectrum that displays
the actual Transmittance spectrum of the plastic
•
Record the spectrum for each color on your report sheet
Then, on your report sheet, explain the correspondence between the color of
light that is transmitted through an object, and the transmittance spectrum.
25
Experiment 2. The relationship between transmittance and the amount of absorbing
material
As you might expect, if you increase the amount of absorbing material, the amount
of light that gets through the material will decrease.
•
Choose any one color, and cut five more sheets of plastic the same size as
your first sheet.
•
Obtain a new spectrum for a single sheet of plastic, as you did in
experiment 1.
•
Locate the table of transmittance values (left side of screen). Move the
slider for this table up and down until you locate a wavelength that has a %
transmittance of about 60%. Record this wavelength on your report
sheet.
•
The math will be a whole lot easier if we change from % Transmittance to
simply T, or fraction of light getting through the sample. Do this by
dividing your % T numbers by 100% so your get a simple fractional
number.
•
On your reportsheet predict the fraction of light that will be
transmitted when you have 1, 2, 3, 4, and 5 sheets of plastic in the light
beam. Have your instructor initial your predictions before you go any
further.
•
Now do the experiment. Record the %transmittance values for 0, 1, 2,
3, 4, and 5 sheets of plastic at your chosen wavelength, then calculate
the transmittance of light (T) using the equation T = %T/100%
•
Finally make a rough plot of this data on your report sheet.
Examine this final plot. Try to come up with an equation that you can use to
predict Ttotal from the number sheets with a T of T1, and write this equation on your
report sheet. Make sure you compare your equation to the real data to see if the equation
and the data fit with one another. Feel free to consult with your instructor if you can’t
come up with a workable equation.
Your equation should have a number raised to a power in it, so it is an exponential
function. That is why your %T data is curved, because it follows an exponential curve.
Scientists are always looking for ways to transform non-linear, curved data, into
nice simple linear functions so they can make a nice linear predictions from the ‘line of
best fit’. Thinking back to algebra, how do you get rid of exponential functions and make
them linear?
Hopefully you remember that taking a log of an exponential function gets rid of
the exponent. Do that now: Take the log of T and plot that against the number of
sheets.
At this point you should have a nice linear plot. The only drawback is that the Y
values have negative numbers, and most people like to work with positive numbers. How
26
do we turn negative number into positive numbers? By multiplying by a negative 1. Do
this to your data and plot the data one more time. Thus the complete transformation
to change our %T data into a nice positive data that is linear and directly proportional to
the amount of material is -log(%T/100%) -or- -log(T)
This mathematical transformation, -log(T), is called Absorbance. While it sounds
like it will be a real pain to calculate -log(%T/100%) for every data point, the nice thing
about the computer is that you can tell the computer to do this calculation for you, so you
never have to do by had again. Let’s set the spectrophotometer to record absorbance
instead of %T.
Click ‘Experiment’-‘Change Units’-Spectrometer 1’ - ‘Absorbance’
Note: if nothing happens when you click on ‘Experiment’ You have to tell the
computer to stop taking data by clicking on the red stop button at the top of the screen.
So now we have discovered that -log(T) = Absorbance; and Absorbance is directly
proportional to the amount of light absorbing material in the light beam, and this can be
stated mathematically as:
A = -log(T) = K × light absorbing molecules.
Where K is some proportionality constant that could be found from the slope of
our line of best fit.
Notice the inverse relationship between absorbance and transmittance. When light
is passes through a material and is not absorbed, the T or %T is high, and the A or
absorptivity is low. On the other hand when light is absorbed by a material, little light
gets through the material so the T or %T is low, but the absorptivity is high.
In Experiment 2 you increased the amount of light absorbing molecules by putting
more sheets of plastic in the light beam. This corresponds to making the amount of
plastic that the light had to go through thicker and thicker. Thus we can restate the above
equation as:
A = K × thickness of material
The term ‘thickness of material’ is called ‘pathlength’ and often abbreviated ‘R’.
So:
A=K×R
If the light absorbing material is a chemical dissolved in a solute, what is another
way we can increase the number of light absorbing molecules in the light beam?
Of course, we can increase the concentration of the absorbing molecules in the
27
solute. At this point you could do another set of experiments exactly the same as the last,
only instead of increasing the number of light absorbing molecules by adding more strips
of plastic, you simply increase the concentration of molecules in the solution. Since the
results are exactly the same, I won’t make you do this again. Instead I will cut directly to
the final equation:
A = K’ × concentration
or
A = K’ × c
Just as we got the ideal gas law equation by combining equations for the individual
gas laws, we can combine our two absorbance equations into one more general equation:
A = K” × R × c
This is called Beer’s law. Just in case you want to know the connection with Beer.
“The law was discovered by Pierre Bouguer before 1729. It is often mis-attributed to
Johann Heinrich Lambert, who cited Bouguer's Essai d'Optique sur la Gradation de la
Lumiere (Claude Jombert, Paris, 1729) — and even quoted from it — in his Photometria
in 1760. Much later, August Beer extended the exponential absorption law in 1852 to
include the concentration of solutions in the absorption coefficient.” (Wikipedia - ‘BeerLambert Law’, Jan 17, 2011.)
In chemistry for aqueous solutions the pathlength is usually given in cm, and the
concentration is usually given in M. In this case the K is called Molar Absorptivity and
the equation is given as:
A=å×R×M
Since A is a unitless number, the constant (å) and has units of L/(mole@cm) to make
the units in the above equation cancel. (Note: å values will change at different wavelengths
just as the Absorbance values change at different wavelengths. If you want to use this
equation to calculate the concentration of an unknown sample, you must use the same
wavelength to measure absorbance in both the standard, which you know the
concentration of , to calculate å; and in the unknown sample.)
28
Experiment 3. Exploring Beer’s Law in a Two-Component System
•
•
•
•
•
•
•
•
•
•
•
•
•
•
•
•
•
•
Make sure your spectrometer is switched over to measure absorbance.
Since the spectrometer has been sitting, run the calibration again. This time,
since we will be running a sample in a cuvette that has water in it, when the
calibration program asks to ‘place a blank cuvette in the device, insert a
cuvette that contains only water. Note: Place the spectrophotomer so that
you can read the name on it left to right, then - If you are using a
SpectroVis Plus spectrometer, the cuvette should be oriented so the clear
sides face the right and left sides of the sample holder. If you are using a
Red Tide spectrometer, the cuvette should be oriented so the clear sides
face front and back in the sample holder.
Remove the cuvette from the spectrometer and rinse it three times with a
blue dye solution.
Obtain the absorbance spectrum of the Blue Dye solution.
Sketch this spectrum on the report sheet
Obtain an exact absorbance value for the blue dye at the wavelength
with the maximum absorbance, and record these values (Wavelength
and absorbance) on your report sheet.
Given that the dye concentration is .001M and the cuvette has a 1 cm
pathlength, determine the å of blue dye at this wavelength.
Remove the cuvette from the spectrometer and rinse it 3 times with the
yellow dye solution.
Obtain the absorbance spectrum of the Yellow Dye solution.
Sketch this spectrum on the report sheet
Obtain an exact absorbance value for the yellow dye at the wavelength
with the maximum absorbance, and record these values (wavelength
and aborbance) on your report sheet.
Given that the dye concentration is .001M and the cuvette has a 1 cm
pathlength, determine the å of yellow dye at this wavelength.
Each person in your group should now obtain their own unknown.
Each person should rinse the cuvette three times with the unknown and
obtain the absorbance spectrum of their unknown.
Sketch this spectrum on the report sheet.
Explain the relationship of the color of the unknown to the absorption
spectrum of the unknown. ie: what color is the sample? Explain this by
observing the absorbance spectrum.
Explain the relationship between the unknown and the blue and yellow
dye solutions by comparing the absorbance spectra of the blue and the
yellow standard dye solutions to the spectrum of the unknown.
Obtain exact absorbance values for your unknown sample at the two
29
•
wavelengths that you used to calculate the å values for the yellow dye
and blue dye, and record these values on your report sheet.
Use the absorbance values at these wavelengths, and the å values
derived above to determine the concentration of the blue and yellow
dyes in your unknown solution.
This last problem is actually a bit of an oversimplification. Since the blue dye has a
slight absorbance at the yellow dye’s maximum, and vice-versa, this is actually a two
component system that needs to be solved with two equations and two unknowns. If you
are familiar with handling two equations and two unknowns from a math class, go ahead
and try it. If you don’t know how to do this, I won’t make you worry about this
complication until you get to Analytical Chemistry!
30
Name: _______________________
_______________________
Report Sheet
Spectrophotometery
Experiment 1. Transmittance of Light and color
1. What is the wavelength range of
Blue light
________ (nm)
Green Light
________
Yellow Light
________
Red light
________
Prediction 1. Transmittance of Yellow Plastic
(Initialed by Instructor)
Prediction 2. Transmittance of Red Plastic
(Initialed by Instructor)
31
Prediction 3. Transmittance of Blue Plastic
(Initialed by Instructor)
Actual Results
Transmittance Spectrum of Yellow Plastic
Transmittance spectrum of Red Plastic
32
Transmittance spectrum of Blue Plastic
Explain the correspondence between the color of light that passes through an object, and
the peaks in the transmittance spectrum of the object.
Experiment 2. The relationship between transmittance and the amount of absorbing
material
Color of plastic you are using for this experiment?
Wavelength at which you have 60% transmittance? _______
Prediction 4. - Fill in the following table:
Number of sheets
T = %T/100%
0
1.0
1
.6
2
3
4
5
Instructor initials: ______________
33
Using the same wavelength and color plastic that you used for your predictions, do the
experiment.
Experimental Values - Fill in the table
Number of sheets
T = %T/100%
0
1
2
3
4
5
Plot of Transmittance vs amount of material
Equation that relates Ttotal to Number of sheets of plastic with a given T1:
34
Plot of log(T) vs. amount of material
Plot of -log(T) vs. amount of material
35
Experiment 3. Exploring Beer’s Law in a Two-Component System
Sketch of Blue Dye spectrum
Wavelength of maximum absorption of blue dye: ________________
Absorbance of Blue dye at this wavelength: ___________________
å of Blue dye at this wavelength: _________________
Show how you calculated å
Sketch of Yellow Dye Spectrum
Wavelength of maximum absorption of yellow dye: ________________
Absorbance of Yellow dye at this wavelength: ___________________
å of Yellow dye at this wavelength: _________________
Show how you calculated å
36
Unknown data
Name of person doing unknown: _____________
Number of unknown: ___________________
Sketch of Unknown Spectrum
1. Explain the relationship between the color of the unknown and the absorbance spectrum above.
2. Explain the relationship between the unknown and the blue and yellow dyes.
3. What are the absorbances values for the unknown at the wavelengths you used to
calculate the å values for the yellow dye and blue dye?
Blue Max?
Wavelength
Yellow Max?Wavelength
nm
nm
Absorbance
Absorbance
4. What are the concentrations of blue and yellow dye in your unknown solution? Show
work
Blue
Yellow
37
Experiment 5.
Identification of Unknown Solutions
Purpose: In this experiment, you will learn:
i
how physical and chemical properties can be used to identify substances
i
how deductive logic is used in the lab
i
how litmus paper and other simple chemical tests work to identify
substances
Background
Many substances have unique chemical and/or physical properties which allow the
chemist to identify the substance by simple observations. Most students can recognize
solid KMnO4, elemental copper and elemental gold by sight, gaseous H2S and ammonia by
odor, and nitric acid by its unique reaction with elemental copper. Many other substances
require a sequence of tests and deductive confirmation or elimination to identify.
In this experiment, you will be provided with twelve aqueous solutions, all of which are
clear and colorless, which you will identify through a series of chemical and physical
analyses. The following is a list of the twelve solutions:
HNO3 (aq)
NaOH (aq)
NH3 (aq)
NaCl (aq)
NaC2H3O2(aq)
H2SO4(aq)
LiOH (aq)
K2SO4 (aq)
NH4Cl (aq)
HC2H3O2 (aq)
HCl (aq)
Ba(OH)2 (aq)
Note that four of the compounds are acids, four are bases, and four are salts. Identify
which are which from the above list before coming to the lab. Also note that three contain
chloride ion, two sulfate ion, and two have distinctive odors. Identify these and record
them in your notebook before coming to the lab.
In the lab, you will obtain portions of each of the unknown solutions, classify them as
acids, bases, or neutrals, and then run further tests to identify each unique compound.
Follow your experiment carefully, with an eye on the list of possible compounds. As it
sometimes happens, a compound may be identified by the fact that it did not react
positively to any of the test procedures.
38
At the end of the experiment, you will be asked to fill out a report sheet, identifying each
of the twelve solutions, and presenting information on how you deduced the identification.
Before coming to the lab, review the chemistry of the next section, and plan a strategy for
solving this twelve-solution problem on the most efficient sequence of tests. Prepare your
notebook by setting up tables and gathering data.
Chemical analyses
Litmus paper: The litmus molecule has two distinct colors, depending on the acid-base
condition of the molecule. The acid form is red and the base form is blue. Paper which is
impregnated with one form can be used to identify compounds which have the opposite
acid-base character. Thus, blue litmus paper turns red when it is wetted with an acid,
while red litmus turns blue when wetted with a base. Neutral solutions do not change
either color of paper. Note that a solution which does not change the color of blue litmus
paper may be either basic or neutral, but it cannot be acidic. Only when a change occurs
can a positive identification be made.
BaCl2 (aq): An aqueous solution of barium ion will react with sulfate ion in solution to
form a white precipitate.
Ba2+ (aq) + SO42-(aq) 6 BaSO4 (s)
This reaction can be used to determine the presence of sulfate ion in neutral and acidic
solutions, and can be used in a reverse fashion by the addition of sulfuric acid to a solution
suspected of containing barium ion.
AgNO3 (aq): An aqueous solution of silver ion will react with chloride ion to form a white
precipitate.
Ag+ (aq) + Cl- (aq) 6 AgCl(s)
To a lesser extent it will also react with sulfate ion to also form a white precipitate
2Ag+ (aq) + SO42- (aq) 6 Ag2SO4(s)
So it should NOT be used with solutions that you have already identified as containing
sulfate. It will also react with bases to form solid AgOH, so this test should be used only
with acids and neutral salt solutions.
Physical analyses
Color and odor can be used to identify some unknown solutions, although in this
experiment, all solutions are colorless. Odor, however, can be used for several of the
compounds. As an example, acetic acid is recognizable by its characteristic vinegar odor.
Conversely sodium acetate, a salt prepared from acetic acid, will not have that same odor
unless the acetate ion is converted back to acetic acid by a neutralization process.
39
C2H3O2- (aq) + H+ (aq) 6 HC2H3O2 (aq)
In a like manner, ammonium ion can be converted to the recognizable ammonia.
NH4+ (aq) + OH- (aq) 6 NH3 (aq) + H2O (l)
Procedure
The twelve solutions are located in squeeze bottles along the side bench of the lab and
each is labeled with an identifying number. The other solutions needed for the tests are
located in dropper bottles in the hood. Your instructor will demonstrate the proper use of
litmus paper and micro test tubes, as well as techniques for testing the odor of a
compound.
Place twelve test tubes (standard and micro can be used) in a test tube rack and obtain
about a milliliter of each of the unknowns (do not measure the volume each time - measure
once then estimate). Be sure to label each test tube with the identifying number of the
solution placed in the tube. Test each solution with litmus paper to classify the compounds
as acidic, basic or neutral. Record your results.
Check for identifying odors for the acids and for the bases. This will provide confirming
evidence, identifying two of the unknowns. Record your results.
Continue to work with the remaining acids. Test them for chloride or sulfate ions, and
confirm the presence of two more of the acids. After a test solution has been added to
one of the unknowns, the mixture should be poured out, the test tube rinsed with
laboratory water from your wash bottle, and a new one milliliter aliquot of unknown
obtained for further testing. Do not, however, waste the unknown solutions. Identify
the fourth acid (HOW?), and record your results.
Identify each of the four neutral salt solutions by determining which one contains sulfate
ion, which chloride ion, which ammonium ion and which acetate ion. See the initial
discussions in this handout for the chemical analyses.
Finally, identify each of the remaining basic solutions (HOW?) Check with your instructor
if you have problems and need some ideas.
40
Name:__________________
Report Sheet
Identification of Unknown Solutions
Identify which bottle number corresponds to each of the following unknown compounds.
Also indicate the results of the litmus test for each compound (turned red, blue, neither),
and write a description (comment or chemical equation) of how you deduced the identity
of the compound.
Compound
Number
Litmus
Description
HNO3(aq)
______
_______
___________________________
NaOH(aq)
______
_______
___________________________
NH3(aq)
______
_______
___________________________
NaCl(aq)
______
_______
___________________________
NaC2H3O2(aq)
______
_______
___________________________
H2SO4 (aq)
______
_______
___________________________
LiOH (aq)
______
_______
___________________________
K2SO4 (aq)
______
_______
___________________________
NH4Cl (aq)
______
_______
___________________________
HC2H3O2 (aq)
______
_______
___________________________
HCl (aq)
______
_______
___________________________
Ba(OH)2 (aq)
______
_______
___________________________
41
Experiment 6.
Molecular Models
Purpose
In this experiment you will visualize molecular structures in terms of VSEPR theory by
building models of various molecules and ions.
Background
The building of molecular models can be very beneficial, and can enable you to visualize
molecular structures and reactions in terms of bonding theories and structure concept such
as VSEPR, Valence Bond Theory, and Molecular Orbital Theory. The construction of
models also provides a better understanding of the three-dimensional characteristics of
molecules and may give insight into the relationship between structure and bonding, and
between structure and reactivity.
In order to build a molecular model, one must first determine the expected geometry and
hybridization at each atom using VSEPR Theory and Valence Bond Theory. The
appropriate components (i.e. atom centers, orbitals, sigma bonds, pi bonds, etc.) are
chosen to construct the molecule in question. Since we simply sticking toothpicks into
Styrofoam balls, you will have to pay attention to try to get your bond angles
approximately right.
Procedure
Part I
Identification of geometrical shapes around atomic centers
To show that you have a clear concept of how to build the different geometries, take five
Styrofoam balls and push in toothpicks to construct one model of each of the following
shapes: linear, trigonal planar, tetrahedral, trigonal bipyramidal and octahedral. Show
these to the instructor for approval before going on to Part II.
42
Part II
Molecular shapes and specific molecules
Follow this procedure for each of the molecules or ions listed at the end of this section.
1.
2.
Draw the Lewis electron dot formula. Draw the resonance forms if they are needed.
Determine the number of regions of electron density around the central atom.
(Note: a region of electron density may be a lone electron, a lone electron pair, a
bond electron pair, or a multiple bond.)
3.
Assign the geometric configuration associated with the number of regions of
electron density according to the following table:
Number of regions of
electron density
geometric configuration
bond angles
2
linear
180 o
3
trigonal planar
120o
4
tetrahedral
5
6
trigonal bipyramidal
octahedral
109.5o
axial-axial 180o
axial-equatorial 90o
equatorial-eq. 120o
90o
4.
Arrange the regions of lone and bonded electron density so as to minimize the
magnitude of repulsive energy using the following basic premises:
a. magnitude of repulsions: lone-lone » lone-bond > bond-bond
b. if the bond angle between regions of electron density is greater than 90o,
repulsions are negligible.
5.
On the basis of the positions of the bonded regions of electron density, assign the
molecular structure.
6.
Determine the appropriate hybridization of the central atom or atoms in order to
give the proper geometric configuration of electron density around these atoms.
7.
Choose the appropriate atom center for the central atom or atoms which will
43
produce the desired electronic geometric configuration. Attach the sigma bond and
other orbital parts as needed.
Atoms other than hydrogen which are attached to these central atoms should be
represented by other black balls. If enough extra orbital attachments are available,
use these to represent the non bonded lone pairs on the outlying atoms.
8.
Write down descriptive information such as the molecular shape, the number of
lone pairs of electrons on the central atoms, the approximate bond angles, where
double or triple bonds exist between atoms, where delocalized or resonance
structures exist, whether the molecule is polar or non-polar.
9.
After a model has been constructed, the instructor will check the model, ask
appropriate questions, and then initial your report sheet. In some cases, you can
construct more than one model at a time, or make minor modifications to a model
to show a sequence of structures to the instructor while he is at your station.
In each lab section, four molecules will be chosen for detailed answers on the
report sheet. These will be announced at lab time. However, you ought to work
out the same answers in your own lab notebook for all the others you make. We
will ask the same questions as are listed on the report sheet for all models you
make.
The molecules or ions to be constructed are
A. BeCl2
G. PCl5
M. N2H4
S. HCN
B. BCl3
H. BrF3
N. C2H6
T. CO
C. CH4
I. SF4
O. C2H4
U. CO2
D. NH3
J. SF6
P. C2H2
V. SO2
E. H2O
K. XeF2
F. PCl3
L. XeF4
Q. H2CO
W. SO3
( C is central)
R. CH2F2
X. CO32-
44
General Guidelines for Drawing Lewis Structures
1. Lewis structures are only useful for covalent bonding. Don’t use for ionic
compounds
2. Count the total number of valence electrons all the atoms of the molecules. Do
not count core electrons. Also remember to add electrons if the molecules is an
anion or subtract electrons if it is a cation.
3. Arrange atoms symmetrically, with the least electronegative atom in the
center and the more electronegative atoms around the outside. F and H should
always be on the outside with a single bond. Cl, Br, and I are also usually around
the outside with single bonds.
4. C, N, O, and F, obey the octet rule. C, N, and O are frequently involved in
double or triple bonds.
5. In radicals (molecules with an odd number of electrons) use pairs of electrons
for each bond, then place the unpaired electron on the least electronegative atom
or in a multiple bond.
6. If a single arrangement of atoms can have double bond(s) in more than two
equivalent locations you have resonance forms present.
7. Nonmetals from the third or higher period may have an expanded octet, if they
are the central atom in a structure. Place extra lone pairs on this atom. In cases
of multiple possible structures the one with the lowest formal charge is best.
8. Be, B, and Al may form compounds with incomplete octets. In this case they
need to be the central atom.
9. Begin to recognize groups of atoms that are frequently bonded together in
different structures. Use these as building blocks to speed up your solution to any
complicated structure
10. When finished check your final structure by double checking the total number
of valence electrons
45
Summary of VSEPR Orbital and Molecular Geometries
# electron
regions
Hybridization
Non-bonding
Pairs
Orbital
Geometry
Bond
Angles
Molecular
Geometry
2
sp
0
linear
180
Linear
3
sp2
0
1
trigonal-planar
trigonal-planar
120
<120
Trigonal planar
V-shaped
4
sp3
0
1
2
tetrahedral
tetrahedral
tetrahedral
109
<109
<109
Tetrahedral
Trigonal pyramid
V-shaped
5
dsp3
0
1
2
3
120&90
<120&90
90
180
Trigonal bipyramid
See-saw
T-shaped
Linear
6
d2sp3
0
1
2
90
<90
90
Octahedral
Square pyramid
Square planar
trigonal bipyramid
trigonal bipyramid
trigonal bipyramid
trigonal bipyramid
octahedral
octahedral
octahedral
46
Names:
Report Sheet
Molecular Models
Have your instructor initial next to each model listed below to verify that you have
constructed it correctly. Your instructor will indicate which four of the models you
should consider for the detailed questions on the next page.
I. Geometrical Shapes
1. Linear______________
2. Trigonal
planar_______________
4. Trigonal
bipyramid____________
3. Tetrahedral __________
5. Octahedral __________
II. Molecules
1. BeCl2____________________
2. BCl3 _____________
3. CH4 _____________
4. NH3 _____________
5. H2O _____________
6. PCl3_______________________
7. PCl5 _____________
8. BrF3______________
9. SF4 _____________
10. SF6 ______________
11. XeF2 ____________
12. XeF4____________
13. N2H4 ___________
14. C2H6 ____________
15. C2H4 ____________
16. C2H2 ____________
17. H2CO ____________
18. CH2F2 ___________
19. HCN _____________
20. CO _____________
21. CO2______________
22. SO2_____________
23. SO3 ______________
24. CO32- ___________
47
MOLECULES
Chemical
Formula
Lewis dot
structure
# of regions of
e- density
around central
atom
Geometrical
arrangement of
e- regions
Drawing of
arrangement of
e- regions
shape of the
molecule
hybridization
of the central
atom
drawing of the
molecule with
bond types and
angles labeled
polar or nonpolar
molecule?
48
Experiment 7.
Isomers, Hybridization, and Molecular Orbitals
Purpose: In this experiment you will:
build models of molecules
examine different isomers of molecules with the same molecular formula
relate atom hybridization to molecular structure
investigate computerized models of molecular orbitals
Background
In last week’s lab we built models of molecules and learned about different molecular shapes
that occur as a result of the different arrangements of atoms and lone pairs. This week we will
build models to expand on the shapes we learned last week. We will see that as molecules
become more complex, multiple spatial arrangements of atoms become possible. Molecules that
have the same molecular formula but different molecular shapes are called isomers. We shall
build models of molecules and learn to recognize different types of isomers.
Again, you will use Lewis structures to determine the number of electron regions, and from this,
the geometrical arrangement of electron pairs as well as the molecular shape. This week we will
also identify what type of hybrid orbitals are used to form covalent bonds. A summary of
Valence Shell Electron Pair Repulsion (VSEPR) theory is provided in the table on the following
page.
For this lab, we will use models that already have defined shapes. Based on the number of
electron regions around an atom, you will need to choose the correct model piece. The pieces
with small pegs represent atoms and the tubes will connect atoms, representing covalent bonds.
Remember, single bonds rotate freely, but double bonds do not. Double bonds cannot rotate
because the p orbitals that overlap to form ð bonds must be oriented parallel to each other.
Different types of isomer exist. Once you have drawn Lewis structures and built models, use the
flow chart on the next page to determine the type of isomers you have modeled.
49
50
Procedure:
For stations 1-4:
1.
2.
3.
4.
5.
6.
7.
8.
First draw a Lewis structure, or if possible, multiple Lewis structures.
Determine the number of electron regions around the central atom(s).
From the number of electron regions, determine the geometry of the electron regions and
the hybridization of the central atom(s).
Build all possible molecular models. These models use atoms that have pegs pointed in
the different orientations, so you will need to consider the hybridization of the central
atom(s) and choose the correct piece from the model. Use tubes to make covalent bonds
connecting atoms. You may choose any available color to represent the various atoms,
but use different colors for different elements.
Using the isomer flow chart, determine what type of isomers you have. There may be
just two isomers or up to five isomers possible, depending on the molecular formula. It
may be possible to have more than one type of isomer for a given molecular formula.
Draw each isomer, using hash and wedge bonds to show the three dimensional structure.
Determine whether each isomer is polar or non-polar. If the molecule is polar, draw the
net dipole moment (remember than we draw dipole moments in the direction of electron
flow—pointing toward the more electronegative atoms).
Have each set of isomers checked and signed off by an instructor.
Before moving to the next station, please take apart the models.
For station 5:
Models of various molecules have been built for you. Compare isomers and match the models to
the structures drawn. Answer questions found in this section.
Please do not take these structures apart.
For station 6:
A computer model of the molecular orbitals can be found on the laptops at station 6. Examine
the computer model and answer questions.
51
Include the following information in your notebook (this part doesn’t need to be turned in).
Station 1.
C5H12
Draw all possible Lewis Structures
How many electron regions does each central
atom have?
What types of isomers are there?
Geometrical arrangements of electron
regions:
Label each isomer as polar or nonpolar
Hybridization of central atoms:
If polar, draw the net dipole moment on the
molecule.
C2H2Cl2
Draw all possible Lewis Structures
How many electron regions does each central
atom have?
What types of isomers are there?
Geometrical arrangements of electron
regions:
Label each isomer as polar or nonpolar
Hybridization of central atoms:
If polar, draw the net dipole moment on the
molecule.
52
Station 2.
SeF2Cl4
Draw the Lewis Structure
Draw all possible isomers
How many electron regions does each central
atom have?
Geometrical arrangements of electron
regions:
What types of isomers are there?
Molecular shape:
Label each isomer as polar or nonpolar
Hybridization of central atoms:
If polar, draw the net dipole moment on the
molecule.
SeF3Cl3
Draw the Lewis Structure
Draw all possible isomers
How many electron regions does each central
atom have?
Geometrical arrangements of electron
regions:
What types of isomers are there?
Molecular shape:
Label each isomer as polar or nonpolar
Hybridization of central atoms:
If polar, draw the net dipole moment on the
molecule.
53
PF3Cl2
Draw the Lewis Structure
Draw all possible isomers
How many electron regions does each
central atom have?
Geometrical arrangements of electron
regions:
What types of isomers are there?
Molecular shape:
Label each isomer as polar or nonpolar
Hybridization of central atoms:
If polar, draw the net dipole moment on
the molecule.
XeF2Cl2
Draw the Lewis Structure
Draw all possible isomers
How many electron regions does each
central atom have?
Geometrical arrangements of electron
regions:
What types of isomers are there?
Molecular shape:
Label each isomer as polar or nonpolar
Hybridization of central atoms:
If polar, draw the net dipole moment on
the molecule.
54
Station 3.
C6H4F2
Draw all possible Lewis Structures.
How many electron regions
does each central atom have?
What type of isomers are these?
Geometrical arrangement of electron regions:
Label each isomer as polar or nonpolar.
Hybridization of the central atoms:
If polar, draw the net dipole moment on the
molecule.
C6H10Cl2
Lewis Structure.
Draw all possible isomers (Just focus on the relative orientation of
bonds on the carbon atoms with chlorine atoms attached.)
How many electron regions
does each central atom have?
Hybridization of the central atoms:
Geometrical arrangement of electron regions:
What type of isomers are these?
55
Station 4.
CHFClBr
Draw the Lewis Structure.
How many electron regions
does the central atom have?
Geometrical arrangement of electron
regions:
Draw all possible isomers.
What type of isomers are these?
Molecular shape:
Label each isomer as polar or nonpolar.
Hybridization of the central atoms:
If polar, draw the net dipole moment on the molecule.
Imagine the paper with colored dots represents a binding site for a specific molecule. How many
of your isomers will ‘fit’ into this site (by matching up the colors)?
56
Station 5.
Structures of optical isomers are provided. Match each molecule to the structure drawn below.
For these models,
· black represents carbon
· white represents hydrogen
· blue represents nitrogen
· red represents oxygen.
Optical isomers are mirror images that are not superimposable. Align each pair of isomers so
that they mirror each other.
For each pair of optical isomers, identify the atom that is responsible for the different spatial
arrangement. This is called a chiral atom (Hint: Chiral atoms have four different groups
attached.) Indicate the chiral atom on each of the structures below. What is the hybridization of
the chiral atom?
For two of these pairs of optical isomers, there are also space-filling models of the molecules.
Match the space filling models to the ball and stick models.
Which type of model do you think is a better representation of the actual shape of the molecules?
In which structure is it easier to identify the chiral atom?
57
Station 6.
Molecular Orbitals
A computerized model of the formaldehyde (H2CO) molecule is open on the computer. This
model shows the shapes of the molecular orbitals.
To view a particular molecular orbital, click on one of the orbitals listed in the lower left window
(listed as 1A, 2A, 3A,…); once the desired orbital is selected, click update. In the other window,
you can see the shape of this molecular orbital. You can drag the mouse and rotate the structure.
Where on the molecule is the second lowest energy (2A) molecular orbital located?
Where on the molecule is the 3A molecular orbital located?
Which molecular orbital most closely represents the ð bond? Which atomic orbitals contribute
the most to this molecular orbital (which atomic orbitals have the highest coefficients)?
Does the 9A molecular orbital appear to be a bonding or an antibonding molecular orbital?
58
Names __________________
__________________
Report Sheet
Isomers, Hybridization, and Molecular Orbitals Laboratory
Station 1
1. C5H12 ___________
2. C2H2Cl2 ___________
Station 2
3. SeF2Cl4 ___________
4. SeF3Cl3 ___________
5. PF3Cl2 ___________
6. XeF2Cl2 ___________
Station 3
7. C6H10Cl2 ___________
8. C6H4F2 ___________
Station 4
9. CHFClBr ___________
Station 5
10. alanine ___________
11. carvone ___________
12. thalidomide ___________
Station 6
13. H2CO ___________
59
Station 1.
Draw all of the isomers of C2H2Cl2. Label each as polar or nonpolar and, if polar, draw the
dipole moment. What type(s) of isomers are these?
Station 2.
Draw all of the isomers of XeF2Cl2. Label each as polar or nonpolar and, if polar, draw the
dipole moment.
Why don’t we observe other isomers of this molecule, such as the one shown
to the right?
Station 3.
Draw all of the isomers of C6H4F2. Label each as polar or nonpolar and, if polar, draw the
dipole moment.
What is the hybridization of the carbon atoms? What does this mean for the overall shape
of the molecule?
60
Station 4.
Draw all possible isomers of CHFClBr. What is the hybridization of the central atom?
What type of isomers are these?
Station 6.
Which molecular orbital most closely represents the ð bond? Draw this molecular orbital.
61
Experiment 8.
Copper Compounds
Purpose:
•
•
•
In this experiment you will:
prepare and make observations on a series of copper-containing compounds
learn how to deduce chemical information from visual observations
write equations describing observed reactions
Background
Most beginning chemistry students have an appreciation for the physical properties of a
metal in its elemental state, but do not have a similar understanding of the many
compounds formed by this metal. Your understanding of the experiment will be enhanced
if, at each step, you ask yourself, and attempt to answer a series of questions concerning
the experiment. For example:
1)
What physical changes are occurring?
2)
What chemical changes do these suggest?
3)
How should we describe these changes in equation form?
During this experiment, you will produce and isolate, in sequence, a series of copper
species:
Elemental copper in a processed form (starting material)
Copper (II) nitrate in aqueous solution (equation A)
Copper (II) hydroxide (equation B)
Copper (II) oxide (equation C)
Copper (II) sulfate in aqueous solution (equation D)
Elemental copper in a non-processed form (equation E)
Copper (I) iodide (equation G)
62
At each stage of the sequence, you will set aside a small portion of the reaction mixture
for later comparison, and then carry the remaining mixture on through the next
conversion. At the end of the experiment, you will compare the seven substances, and
write net ionic equations describing each step of the sequence.
The following formula unit equations describe the procedure:
A) Cu (s) + 4 HNO3 (aq) Y Cu(NO3)2 (aq) + 2 NO2 (g) + 2 H2O (1)
B) Cu(NO3)2 (aq) + 2 NaOH (aq) Y Cu(OH)2 (s) + 2 NaNO3 (aq)
C) Cu(OH)2 (s) Yheat CuO (s) + H2O (g or l)
D) CuO (s) + H2SO4 (aq) Y CuSO4 (aq) + H2O (1)
E) CuSO4 (aq) + Zn (s) Y Cu (s) + ZnSO4 (aq)
F) Zn (s) + 2 HCl (aq) Y ZnC12 (aq) + H2 (g)
G) 2 CuSO4 (aq) + 4 KI(aq) Y 2 CuI (s) + I2(s) + 2 K2SO4(aq)
H) I2 (s) + NaHSO3 (aq) + H2O (1) Y NaHSO4 (aq) + 2 HI (aq)
Procedure
The procedure for this experiment is lengthy and complicated, so make sure that you
make careful notations of your observations. Also, label each of the products so that you
can identify each at the end of the period.
In preparation for the work, clean your six standard test tubes, two 250 mL beakers, two
stirring rods, filter flask, Buchner funnel, and graduated cylinders. Obtain a sample of
copper from your instructor.
Experiment 1
Preliminary Observations of Elemental Copper
Break off a small piece of the copper (the length of a grain of rice) and place it on a watch
glass. Record your notes on its appearance, shape, color, etc. Place 0.35 to 0.45 g of the
sample in a 250 mL beaker.
63
Experiment 2
Copper (II) Nitrate
Note: carry out all parts of this step in the fume hood. Add 20 mL of 6 M HNO3 (aq)
to the 250 mL beaker containing your elemental copper sample and cover the beaker with
a plain watch glass. Allow the reaction to continue until all of the elemental copper has
been converted to Cu2+ (aq) (at least 10 minutes). If any unreacted copper metal remains
in the beaker after this time, notify your instructor. Remove the watch glass and carefully
swirl the solution to expel the NO2 (g) that is formed during the reaction and then add
about 25 mL of distilled water to the reaction mixture. The beaker may now be removed
from the hood and taken to your desk.
Pour about 3-5 mL of the solution into a standard test tube. Make notes on its appearance
and continue with the experiment.
Experiment 3
Copper (II) Hydroxide
Measure out 20 mL of 6 M NaOH (aq) into a graduated cylinder. Slowly add about 12
mL of this base to the acidic copper (II) nitrate solution in the beaker, with constant
stirring. Check to see if the solution is slightly alkaline. This is most easily accomplished
by testing the solution with litmus paper. If the solution tests acidic, add 1 more mL of 6
M NaOH, and check again for acidity with litmus paper. Repeat this process until the
solution is slightly alkaline. Avoid a large excess of base which could result in dissolving
the product. Record all observations of the reaction in your notebook.
Pour a few milliliters of the mixture into a second standard test tube. Make notes on its
appearance, and continue with the experiment
Experiment 4
Copper (II) Oxide
Place the beaker containing the aqueous suspension of copper (II) hydroxide on your hot
plate and set the dial to about 3. Heat the suspension until steam starts to form and then
continue heating, with constant stirring, for about 10 minutes. This will assure total
conversion of the hydroxide to copper (II) oxide, help coagulate the precipitate
(sometimes referred to as digestion) and prevent loss of precipitate due to bumping.
During the heating process, use your stirring rod to scrape any blue particles of precipitate
that remain on the sides of the beaker down into the reaction mixture. When the reaction
is complete, all of the blue color of the mixture should be gone, a brown-black solid
should settle to the bottom of the beaker, and the solution should be clear. Pour off the
64
clear supernatant liquid, and wash the precipitate in the beaker with 50 mL distilled water.
Let the precipitate settle and again pour off the supernatant liquid. Record all observations
of the reaction.
Transfer a small portion of the CuO solid (size of a pea) to one of the test tubes. Make
notes on its appearance, and continue with the experiment.
Experiment 5
Copper (II) Sulfate
Slowly add up to 25 mL of 3 M H2SO4 (aq) to the beaker containing the CuO precipitate
until the precipitate just dissolves. Very carefully stir the solution, if necessary, to assure
complete reaction of the CuO. (You may not need the entire 25 mls.)
Pour 3-5 mL of the solution into another standard test tube. Make notes on its appearance,
and continue with the experiment.
Experiment 6
Elemental Copper
Pour half of the remaining CuSO4 (aq) into the other 250 mL beaker. Label one of them
beaker #1, and the other beaker #2. The first will be used to generate elemental copper,
while the second will be used to produce CuI.
Weigh out about 1.0-1.2 g of zinc metal powder. While stirring the solution, slowly add
the zinc to beaker #1. Be careful, the solution will get very hot. Continue to stir the
solution until all the Cu2+ (aq) has been converted to copper metal (the solution becomes
colorless), and until the acid attack on the zinc has subsided (hydrogen gas evolution
decreases). Let the metal particles settle to the bottom of the beaker and then decant (pour
off) about three-fourths of the supernatant liquid. Add 10 mL of 6 M HCl (aq) to the
metal in order to dissolve any unreacted zinc. Heat gently, if necessary, to accelerate the
reaction.
After the zinc metal is totally reacted, add 50 mL of distilled water to the reaction beaker
and stir. Set up a vacuum filtration apparatus with a Buchner funnel. Place a piece of filter
paper into the funnel and filter the elemental copper. Rinse with two 10 mL portions of
water. Transfer the filter paper, containing the elemental copper, to a watch glass. Make
notes on its appearance and proceed with the rest of the experiment.
Since this is an apparatus you will use again, make a diagram of vacuum filtration set-up
in your lab notebook.
65
Experiment 7
Copper (I) Iodide
To the contents of beaker #2, add 20 mL of 0.5 M KI (aq). While stirring the solution, add
dropperfuls of 0.5 M NaHSO3 (aq) until all of the brown color is eliminated. Place the
beaker on the hot plate and heat until it is steaming. Continue to heat and stir for five
minutes, remove the beaker from the hot plate, and allow it to cool.
Set up a vacuum filtration apparatus with a Buchner funnel. Place a piece of filter paper
into the funnel and filter the CuI (s). Rinse the reaction beaker with two 10 mL portions
of distilled water, pouring each rinse through the precipitate to rinse it. Draw air through
the precipitate for a few minutes to dry it. Transfer the filter paper and precipitate to a
watch glass. Make notes on its appearance and proceed to the last step.
Experiment 8
Observations and Report Sheet
Record, in your laboratory notebook, the physical condition of each of the copper
containing materials. In particular, record the color and form (crystalline, metallic,
amorphous) of each.
After your observations are complete, dispose of your solid products in the ceramic jar
above the sink and flush your remaining liquids down the sink drain with running water.
For each product, write the appropriate set of net ionic equations to describe the reaction.
For redox reactions, include the balanced half reactions and the overall net ionic
equations.
66
Name:
Report Sheet
Copper Compounds
I. Writing Equations
Each of the following chemical equations describes one of the reactions that you carried
out in this week’s laboratory assignment. For each, write in the space provided the ionic
and net ionic equations that describe the reaction. For reactions A, E, and F also write
the balanced oxidation and reduction half reactions.
A) Experiment 2
Cu (s) + 4 HNO3 (aq) Y Cu(NO3)2 (aq) + 2 NO2 (g) + 2 H2O (l)
B) Experiment 3
Cu(NO3)2 (aq) + 2 NaOH (aq) Y Cu(OH)2 (s) + 2 NaNO3 (aq)
C) Experiment 4
Cu(OH)2 (s) YHeat CuO (s) + H2O (g or l)
67
D) Experiment 5
CuO (s) + H2SO4 (aq) Y CuSO4 (aq) + H2O (l)
E) Experiment 6
CuSO4 (aq) + Zn (s) Y Cu (s) + ZnSO4 (aq)
F) Experiment 6
Zn (s) + 2 HCl (aq) Y ZnCl2 (aq) + H2 (g)
G) Experiment 7
2 CuSO4 (aq) + 4 KI (aq) Y 2 CuI (s) + I2 (s) + 2 K2SO4 (aq)
68
H) Experiment 7
I2 (s) + NaHSO3 (aq) + H2O (l) Y NaHSO4 (aq) + 2 HI (aq)
II. Observations of Physical Properties
Complete the following table by indicating your observations. Indicate whether solid
compounds are metallic, crystalline, or amorphous.
Compound
Color
Physical Form
Cu
Cu2+
NO2
Cu(OH)2
CuO
Zn2+
H2
CuI
69
Experiment 9.
Synthesis of a Cobalt Salt
Purpose
In this experiment you will synthesize an ionic compound that will be used in
laboratory experiments during the next couple of weeks. The name of the compound is
cobalt (II) oxalate dihydrate, and its formula is CoC2O4 . 2 H2O.
Background
Synthesis is a process by which a chemist prepares a compound of interest from other
(usually simpler) elements or compounds. Several steps are normally required, depending
on the complexity of the compound formed and the nature of the starting materials. In this
experiment. the formation of the cobalt compound requires dissolving the starting
materials, mixing them, precipitating the product. recovering the product by vacuum
filtration, and removing impurities in the recovered product. The product will then be
dried and massed to determine the percentage yield for the experiment.
The final reaction today can be described by the following chemical equation:
CoSO4(aq) + (NH4)2C2O4 (aq) + 2 H2O(l) Y CoC2O4 . 2 H2O(s) + (NH4)2SO4 (aq) RXN1
However, the starting materials for this synthesis are two hydrated forms of the cobalt (II)
and oxalate ions: cobalt (II) sulfate heptahydrate, CoSO4 . 7 H2O (s) and oxalic acid
dihydrate. H2C2O4 . 2 H2O (s), and do not appear directly in the above molecular equation.
You actually weigh out CoSO4 . 7 H2O (s) and place this into water where the reaction
shown below occurs:
CoSO4 . 7 H2O (s) 6 CoSO4(aq) + 7 H2O(l)
RXN2
Note that since the starting material in this reaction has 7 molecules of water bound for
every one molecule of CoSO4, you must include the mass of these additional water
molecules in the molar mass of this material.
The oxalate is a bit more complicated. H2C2O4 . 2 H2O (s) itself does not like to dissolve
in water, but it dissolves much better in a mixture of ammonia and water as shown in
equation 3:
H2C2O4 . 2 H2O (s) + 2NH3(aq) 6(NH4)2C2O4(aq) + 2 H2O (l)
RXN3
70
Before coming to the lab, calculate the molar mass of each of the starting materials and
the final product. Record these values in your notebook. These numbers will be needed to
calculate the limiting reagent and theoretical yield from your starting masses.
Procedure
Weigh approximately 5.0 - 5.2 g CoSO4 . 7 H2O (s) on a piece of weighing paper to .001 g
significance and place in another clean, dry 250 mL beaker. Add about 100 mL distilled
water to the beaker and stir until all of the salt is dissolved.
Weigh approximately 2.2 - 2.4 g H2C2O4 . 2 H2O (s) on a piece of weighing paper.
Record your weight to .001 g significance, and place the solid in a clean, dry 250 mL
beaker. Add about 100 mL distilled water and 2 mL of 7 M NH3 (aq) to the beaker. Stir
the mixture until all the acid is dissolved.
Warm the beaker slightly on a hotplate if necessary to dissolve the solid.
Slowly pour the Co2+ (aq) solution, a few mL at a time into the oxalate ion solution with
constant stirring. After all of the cobalt (II) ion solution has been added, place the beaker
containing the mixture in an ice bath to induce and complete the precipitation of the
product. Stir the reaction mixture occasionally to equilibrate the temperature and to effect
the formation of larger more easily filtered particles.
Allow the precipitate to settle in the reaction mixture and set up a vacuum filtration
system with a Buchner funnel. Also prepare about 30 mL of ice cold distilled water by
partially filling a 50 mL graduated cylinder and placing it in the ice bath to cool.
Insert a piece of filter paper into the funnel and turn on the water aspirator vacuum.
Slowly pour the supernatant liquid through the funnel transferring the solid with the last
few milliliters. Rinse the beaker with about 10 mL of ice cold water and pour this over the
precipitate. Repeat with a second 10 mL portion of cold water. This process will both
facilitate total recovery of the precipitate and wash out impurities. Draw air through the
funnel for a few minutes to partially dry the precipitate. Disconnect hose between flask
and water tap BEFORE you turn off the water to turn off the vacuum. Transfer the filter
paper containing the product to a small beaker. Carefully store the labeled beaker in the
hood to allow the product to dry until the next laboratory period.
71
Before leaving the lab record the masses of the starting materials on your report sheet.
Determine the limiting reagent and calculate the theoretical yield. Set up the remaining
calculations on the report sheet. which will be completed once final data is collected
during the next lab period. Check the results of your calculations with the instructor, who
will initial the sheet.
Next Laboratory Period
Remove the CoC2O4 . 2 H2O (s) product from the hood (carefully!!). Weigh a clean, dry
vial on the analytical balance. Transfer the product to the vial and reweigh to determine
the mass of product obtained. Record this final product mass on your report sheet from
last week calculate the percentage yield and turn in the report sheet. Retain the vial of
product since you will use this salt for two more experiments.
[NOTE: At the end of the term, during checkout, you should dispose of any
remaining product in the container provided by the instructor.]
72
Name:
Report Sheet
Synthesis of CoC2O4@2H2O
1. Moles of CoSO4
A. Mass of CoSO4@7H2O used ____________________
B. Moles of CoSO4 produced in RXN 2 _________________
C. Calculation to find moles of CoSO4. (Show equation)
2. Moles of (NH4)2C2O4
A. Mass of H2C2O4@2H2O used ____________________
B. Moles of (NH4)2C2O4 produced in RXN 3 _______________
C. Calculation to find moles of (NH4)2C2O4. (Show equation)
73
3. Limiting reagent
A. Moles of CoC2O4@2H2O produced in RXN 1 if CoSO4@7H2O is limiting reagent?
______________________
B. Show equation for above calculation:
C. Moles of CoC2O4@2H2O produced in RXN 1 if H2C2O4@2H2O is limiting reagent?
________________________
D. Show equation for above calculation:
E. Limiting Reagent? ________________________
4. Theoretical yield
A. Based on the above limiting reagent, what is your theoretical yield of
CoC2O4@2H2O in grams?
____________________________(g)
B. Show equation used to obtain the above value
3. % Yield
A. Yield of CoC2O4@2H2O ___________________(g)
B. % yield for reaction? __________________
C. Show equation for above answer
74
Experiment 10.
Reaction Stoichiometry
Purpose
In this experiment you will determine the mass and percentage of cobalt in
cobalt oxalate dihydrate by thermal decomposition.
Background
Many compounds upon heating decompose into an oxide of known composition. Such is
the case with the cobalt salt you have synthesized.
3CoC2O4 . 2 H2O (s) + 2 O2 (g) Y Co3O4 (s) + 6 CO2 (g) + 6 H2O (g)
In this lab you will start with a known mass of the oxalate salt, heat it until complete
decomposition has taken place, and then determine the mass of the product. From the
formula of the product, you can determine the mass of cobalt in the oxide and the
percentage cobalt in the original oxalate salt.
Procedure
Wash and dry a crucible and lid. Place them on a hotplate set at high and heat for ten
minutes to completely dry. At the end of the heating period remove the crucible and lid
carefully from the hotplate and place on a casserole to cool for ten minutes.
Take the cooled crucible and lid into the balance room and weigh them together to the
nearest 0.001 g. Transfer about 0.3 g of powdered cobalt (II) oxalate dihydrate (your
product in the vial) to the crucible and reweigh accurately to obtain the mass of the salt.
Place the covered crucible containing the salt on the hotplate [set at high], and allow the
crucible to heat for one hour.
75
At the end of the hour, remove the crucible and lid carefully from the hotplate, place it on
the casserole to cool for ten minutes, and then reweigh to obtain the mass of the cobalt
oxide product.
Determine the mass of cobalt in your product from the measured mass of cobalt oxide and
the known molar masses of cobalt and cobalt oxide. Using this and the mass of cobalt
oxalate dihydrate that you started with calculate the (experimental) percentage cobalt in
your cobalt oxalate salt. For comparison, calculate the theoretical percentage of cobalt in
cobalt oxalate dihydrate using only the molar masses of these substances. Do the two
results agree? Why or why not? Transfer your results and sample calculations from your
notebook to your report sheet.
76
Name:
Report Sheet
Reaction Stoichiometry
1. Raw data
A. Mass of salt (reactant)
Mass of Crucible and salt
___________________g
Mass of Crucible alone
- ___________________g
Mass of salt (CoC2O4@2H2O)
B. Mass of Oxide (product)
Mass of Crucible and product
Mass of Crucible alone
_____________________g
___________________g
- ___________________g
Mass of product (Co3O4)
___________________g
2. Data Analysis
A. Mass of Co in oxide product (equals Co in starting reactant) ______________g
B. Calculation for determination of Co in oxide product (or starting reactant):
C. Experimental % Co in CoC2O4@2H2O salt
__________________%
D. Calculation for determination of experimental % Co in CoC2O4@2H2O salt:
77
E. Theoretical % Composition of Co in CoC2O4@2H2O? ___________________%
F. Equation for determination of theoretical % Co in CoC2O4@2H2O salt:
G. Based on how your experimental results compare to the theoretical results,
comment on the purity of the CoC2O4@2H2O salt you analyzed.
78
Name(s)__________________
Chem 112L
Equations Worksheet
1. Write balanced molecular, complete ionic, and net ionic equations for the reaction of Silver
nitrate with Potassium iodide in aqueous solution.
2. Write balanced molecular, complete ionic, and net ionic equations for the reaction of
Copper(II) chloride and sodium hydroxide in aqueous solution.
79
3. Write balanced molecular and complete ionic equations for the reaction of HClO4 and
Mg(OH)2. (Hint: the net ionic reaction is H+(aq) + OH-(aq) 6 H2O (l)
4. Write balanced molecular and complete ionic equations for the reaction of nitric acid and
aluminum hydroxide. (Hint: the net ionic reaction is H+(aq) + OH-(aq) 6 H2O (l)
80
Equations Worksheet - Alternate Second Page
A. Necessary information
The products of the reaction in problem 1 are solid silver iodide, and aqueous potassium
nitrate.
The products of the reaction in problem 2 are solid copper(II) hydroxide, and aqueous sodium
chloride.
3A. If I start the reaction in problem 1 with 42.5 g of silver nitrate and 39.8 g of potassium
iodide, which compound is my limiting reactant?
3B. If the reaction in 3A gives me 55.0 g of silver iodide, what is the % yield for the reaction?
4A. If I start the reaction in problem 2 with 16.0g of copper(II) chloride and 6.00g of sodium
hydroxide, which compound is my limiting reactant?
4B. If the reaction in 4A gives me 9.75g of copper(II) hydroxide product, what is the % yield for
the reaction?
81
Experiment 11.
Redox Titration
Purpose
In this experiment you will use redox titrations to determine the mass
percentage composition of the oxalate ion in the salt cobalt (II) oxalate dihydrate.
Background
In this experiment you will determine the percentage of oxalate ion in the cobalt (II)
oxalate dihydrate you prepared two weeks ago. The analysis will be done by means of a
titration, using a redox reaction. Read pages l61-164 of your text (Zumdahl) to learn
about titrations, standardization, primary and secondary standards, burets, equivalence
and end points and titration calculations. Note that you will be doing an oxidationreduction titration rather than an acid base titration.
You will use sodium oxalate (Na2C2O4) as a primary standard to determine the
concentration of a solution of KMnO4. The latter will then serve as a secondary standard
to determine the percentage oxalate ion in the cobalt (II) oxalate dihydrate by titration. In
these titrations KMnO4 acts as an oxidizing agent, while the oxalate ion acts as a reducing
agent. Potassium permanganate is particularly useful in these titrations because the
strongly colored MnO4- (aq) species yields a colorless product, Mn2+ (aq), so when a
slight excess of the titrant is added, the reaction mixture becomes pink. Thus the
appearance of a permanent pink color signifies the end of the reaction. Potassium
permanganate, then, serves as its own indicator for the end-point of the titration.
The following reactions describe the chemistry involved in this experiment. Potassium
permanganate, sodium oxalate, and cobalt (II) oxalate are all strong electrolytes, and
completely dissociate into their constituent ions in solution:
KMnO4 + H2O Y K+ (aq) + MnO4- (aq)
Na2C2O4 + H2O Y 2 Na+ (aq) + C2O42- (aq)
CoC2O4 + H2O Y Co2+ (aq) + C2O42- (aq)
82
When sulfuric acid is added to solutions of oxalate ions, the hydrogen ions in the sulfuric
acid solution convert the oxalate ion to oxalic acid:
C2O42- (aq) + 2 H+ Y H2C2O4 (aq)
The species involved in the redox reactions undergo the following changes:
MnO4- (aq) + 8 H+ (aq) + 5 e- Y Mn2+ (aq) + 4 H2O (l)
H2C2O4 (aq) Y 2 CO2 (g) + 2 H+ + 2 eThe overall net ionic equation of the titrations, then, is
2 MnO4- (aq) + 5 H2C2O4 (aq) + 6 H+ (aq) Y 2 Mn2+ (aq) + 10 CO2 (g) + 8 H2O (l)
Theoretically, this reaction is too slow to provide a good titrimetric method, but its speed
can be increased by heating the sample to about 600C. Also, addition of a catalyst can
speed up the reaction, and, in this case, one of the products (Mn2+) acts in this capacity. It
is for this latter reason that the first addition of the permanganate ion titrant will require
several seconds to react, but subsequent additions will react quite quickly.
Procedure
Clean and rinse the flask on your desk. Fill it with about 500 mL of distilled water and
add 50 mL of 0.20 M KMnO4 (aq). Put a piece of parafilm over the top and swirl the flask
to completely mix the permanganate reagent. This is your permanganate solution which
will be standardized and then used in the determination of the oxalate ion in the
synthesized cobalt salt. You will have serious errors in your determination if your
permanganate solution is not homogeneous. (i.e. Not completely mixed.)
Put a 250 ml beaker filled with distilled water on your hotplate and heat it until it is
almost too hot to touch. Next, directly weigh on weighing paper, four samples of 0.200 g
to 0.250 g each of Na2C2O4. Record these weights in your notebook, and transfer the
material to clearly labeled Erlenmeyer flasks that have been cleaned and rinsed with
distilled water.
83
Obtain a buret. empty it, rinse it with two small (5 - 10 mL) portions of your diluted
permanganate solution and then fill it. Check for and remove any air bubbles in the tip of
the buret. Record the initial buret reading.
When you are ready to do your first titration pour about 50 mls of hot distilled water into
one of your flasks containing a Na2C2O4 sample. Next add 10 mL of 6 M H2SO4(aq) to
the flask and begin the titration by adding a few drops of the permanganate solution while
swirling the solution in the flask. Wait for the solution to clear before you add any more
permanganate. If necessary, put the flask on the hot plate to warm as you wait for it to
clear. Once it is clear you can continue the titration by slowly adding titrant and swirling
the flask. Continue to add the titrant until the end-point is reached as indicated by the
appearance of a very faint pink color that persists at least 30 seconds. The temperature of
the solution should not drop below 60oC during the titration. Record the final buret
reading and calculate the concentration of the permanganate solution. Check with your
instructor to make sure that the concentration is in the proper range before proceeding.
Titrate the remaining sodium oxalate samples in the same fashion and calculate the
average concentration of your titrant.
Empty the Erlenmeyer flasks and rinse each well with distilled water. From the vial
containing cobalt (II) oxalate dihydrate, weigh out, on the analytical balance four samples
of 0.200 g to 0.250 g each into the Erlenmeyer flasks. When you are ready to titrate the
material in a given flask, add 50 mL of hot distilled water and 10 mL of 6 M H2SO4(aq).
As before start the titration with a few drops of titrant, and wait for the solution to turn
clear (but pink) before titrating to the endpoint. Although the end-point will not be as
distinct as with the sodium oxalate solutions since the cobalt ion will contribute to the
color of the solution an apparent color transition marking the end-point will occur.
Calculate the percentage oxalate ion in each of the titrated samples. Fill in the report
sheet, and give it to the instructor.
84
Name:_________________
Report Sheet
Redox Titration
I. Standardization of Potassium Permanganate Solution
A. Trial I
1. Mass of Na2C2O4 (s) __________________g
2. Moles of Na2C2O4 (s) ___________________ mol
3. Calculation of moles (show set-up):
4. Equivalent moles of KMnO4 ________________ mol
5. Calculation of equivalent moles (show set-up):
6. Volume of KMnO4(aq) ________________mL
7. Molarity of KMnO4(aq)_______________ M
8. Calculation of molarity (Show set-up):
85
B. Other trials
Mass of Na2C2O4
Trial 2
Trial 3
_____________
____________
Moles of Na2C2O4
Volume of KMnO4(aq)
Molarity of KMnO4(aq)
C. Final Result:
Molarity of KMnO4(aq)
±
(Average)
mol/L
(Standard deviation)
1. Calculation of average (show set-up):
2. Calculation of Standard Deviation (show set-up, see Appendix 1.):
II. Determination of Percentage Oxalate in your Salt
A. Trial I
1.Mass of CoC2O4@2H2O (s) ______________g
2. Volume of KMnO4
mL
3.Mass of C2O42- in salt ___________________g
4.Calculation of mass (show set-up):
86
5.Percent C2O42- in salt ______________________
6.Calculation of Percentage (show set-up):
B. Other Trials
Trial 2
Trial 3
1.Mass of CoC2O4 . 2H2O (s) _____________
2. Volume of KMnO4
___________
3. Mass of C2O42- in salt
_____________
4.Percent C2O42- in salt
______________
C. Final Results
1.Percent C2O42- in salt _________________%
(Avg.)
±
(Std. Dev.)
2.Calculation of average (show set-up):
3.Calculation of Standard Deviation (Show set-up):
4. Theoretical % of C2O42- in salt based on molecular mass_____________
87
5. Calculation of Theoretical % C2O42- in salt based on molecular mass:
88
III Summary of Cobalt Synthesis results
1. % yield of CoC2O4@2H2O
_____________________(If synthesized)
2. Actual % of Co in Salt _____________________
3. Theoretical % of Co in Salt _____________________
4. Actual % oxalate in Salt
_____________________
5. Theoretical % Oxalate in Salt_____________________
6. Based on the above data, evaluate your synthesis:
89
Experiment 12.
The Ideal Gas Law
Purpose
In this experiment you will discover how P, V, T, and n are related to each other for an ideal
gas.
Background
Boyle first discovered the relationship between gas pressure and volume in the late 1600's.
Charles then described the relationship between a gas’s volume and temperature around 1800.
Finally in 1811 Avogadro described the relationship between a gas’s volume and the number of
particles in the gas. In today’s three hour experiment you will use modern technology to
discover for yourself these basic principals that took about 200 years to uncover!
Experiment I. The relationship between P and V
Turn on your lab Quest or computer interface and plug in the pressure sensor. Go to the sensors
tab and then find the change units pull down. Set the units of the sensor to atm (atmosphere).
Now find the 20 mL syringe. Pull the plunger on the syringe so it reads 5.0 mL. Now use the
luer lock fitting to attach the syringe to the pressure sensor. Record the pressure and volume.
Now, by moving the plunger in and out, record at least four other pairs of pressure and volume.
Experiment II. The relationship between n and P
Find your 125 mL Erlenmeyer flask. While it says 125 mLs, it actually holds a bit more. You
need to determine the actual total volume of the flask. Fill the flask with water. Now pour that
water into a graduated cylinder so you can measure the total volume of the flask. Record this
volume. The volume of the tubing and fittings is about 4 mL. So what is the total volume of
your flask and the tubing? Dry your flask when you have finished, because it has to be dry for
the remaining experiments.
While you might think you need to know ‘n’, the number of moles of air to do this experiment
you really don’t. All you need is a number that is proportional to n. The number we will use is a
zorkblat. We will assume that a 10 ml unit of air in this room contains 1 zorkblat of molecules.
Calculate the of zorkblats of air contained in your flask and tubing combined.
Find the white stopper with the luer-lock fitting, and press it tightly into a dry 125 mL
Erlenmeyer Flask. Attach one end of the luer-lock tubing to your pressure sensor and the other
end to the luer lock fittings on the flask that does not have a valve on it. Record the initial
pressure. Now fill the syringe with 10 mL of air (1 zorkblat of molecules). Place the syringe on
the luer-lock fitting that has the valve on it. Deliver 1 zorkblat of molecules into the flask, close
the valve and record the pressure. Remove the syringe, fill it with another zorkblat of molecules,
Attach the syringe to the apparatus, open the valve, deliver the molecules to the flask, close the
90
valve and record the pressure. Continue this procedure until you blow out your stopper. (When I
tested this, I blew out the stopper when I put the fifth zorkblat into the flask). For each addition
of air calculate the number of zorkblats using the equation:
Experiment III. The relationship between V and T
Initial set up and first point
Find the white stopper with the luer-lock fitting, and press it tightly into a dry 125 mL
Erlenmeyer Flask. Attach one end of the luer-lock tubing you your pressure sensor and the other
end to one of the luer-lock fittings on the flask. Set the plunger on the syringe to 10 mL and then
attach the syringe to the other luer-lock fitting on the flask. Record the temperature (in Kelvin),
pressure and volume on the syringe.
Low temperature point
Now take the assembly back to the sink that has the large ice bath. Submerge the temperature
sensor, flask, syringe, and most of the tubing in the ice bath, but be careful to keep the pressure
sensor out of the water. After the system has equilibrated for 3-5 minutes, move the plunger on
the syringe up and down until the pressure reads the same as the pressure in the room. Record
the volume of the system (syringe, tubing and flask) and the temperature in the bath.
High temperature point
Now take the assembly back to the sink that has the large warm water bath. Submerge the
temperature sensor, flask, syringe, and most of the tubing in the warm water bath, but be careful
to keep the pressure sensor out of the water. After the system has equilibrated for 3-5 minutes,
move the plunger on the syringe up and down until the pressure reads the same as the pressure in
the room. Record the volume of the system (syringe, tubing and flask) and the temperature in
the bath.
Experiment IV. The relationship between P and T
You will not use the syringe for this experiment, so simply close the valve on the syringe filling
and remove the syringe. Now clamp this into a ring-stand set up. Get a 600 mL beaker and fill it
½ full with ice, then add water until it is about 3/4 filled, and place it on a hot plate (do not turn
it on yet!) Use the ringstand to submerge the flask as much as possible in your water bath and
put the temperature sensor in the water bath as well. Make sure the tubing and all electrical
wires do NOT touch the hotplate. Record the initial pressure and temperature. Double check
that no wires are touching the hotplate, and turn it on high. You can use the temperature sensor
to stir the water in the bath (Watch our for the wire!). Record the pressure and the temperature
about every 5o. Make sure you record your temperature in K! Turn off the heat, remove the
apparatus from the water bath and end your experiment after you have recorded 8-10 pairs of
temperature and pressure.
Make sure you make a diagram of each of the different experimental set-ups in your lab
notebook.
Name:_____________________
91
_____________________
Report Sheet
The Ideal Gas Law
Experiment I.
Raw Data
Volume (L)
Relationship between V and P
Pressure (atm)
Plot the above data on a piece of Graph paper and attach it to this report sheet
What does the graph tell you about the relationship between V and P? _____________________
Write the above relationship in an equation. ___________________________
If the above equation has a constant, determine the value of the constant for your data.
KI =________________
(Don’t forget units!)
Experiment II. The relationship between n and P
Raw Data
Volume of flask and tubing______
Zorkblats of air in flask and tubing ___________
Zorkblats
Pressure (atm)
Plot the above data on a piece of Graph paper and attach it to this report sheet
92
What does the graph tell you about the relationship between n and P? _____________________
Write the above relationship in an equation. ___________________________
If the above equation has a constant, determine the value of the constant for your data.
KII =________________
(Don’t forget units!)
Experiment III. The relationship between T and V
Raw Data
Volume of flask_____________
Pressure in system _____________
Temperature (K)
Volume (L)
Plot the above data on a piece of Graph paper and attach it to this report sheet
What does the graph tell you about the relationship between T and V? _____________________
Write the above relationship in an equation. ___________________________
If the above equation has a constant, determine the value of the constant for your data.
KIII =________________
(Don’t forget units!)
93
Experiment IV. The relationship between T and P
Raw Data
Temperature (K)
Pressure (atm)
Plot the above data on a piece of Graph paper and attach it to this report sheet
What does the graph tell you about the relationship between T and P? _____________________
Write the above relationship in an equation. ___________________________
If the above equation has a constant, determine the value of the constant for your data.
KIV =________________
(Don’t forget units!)
Now, put all this together into one equation
P=...............
If the above equation has a constant, determine the value of the constant from your data.
Kcomplete =________________
Express K in units of atm, K, L and Zorkblats
Show how you calculated K:
Don’t forget to attach all your graphs!
94
95
Experiment 13.
Molar Mass of a Vapor
Purpose: In this experiment, you will:
•
measure the mass, temperature, pressure, and volume of an
unknown vapor
•
use the ideal gas law to calculate the molar mass and
•
predict the identity of your sample
Background
The molar mass of a compound is one of its most fundamental properties. When a new
compound is synthesized or identified, its molar mass is usually one of the first properties
to be determined. A number of different methods can be used to determine molar mass,
depending on the properties of the compound. For volatile liquids, molecular substances
having relatively low boiling points, the molar mass can be determined by measuring the
mass of the vapor in a fixed-volume container at a known temperature and pressure.
As shown in equation 5.1 from your text, page 204, the molar mass of a compound is
directly related to its density.
(1)
Where d is the density of the gas in g/l, R is the gas constant (0.082059 L@atm/K@mol), T
is the absolute temperature and P is the pressure in atm. In this experiment you will
measure the volume of a container and the mass of a gas in the container, so the above
density term can be replaced with:
(2)
Where m is the mass of the gas measured in grams and V is the volume of the container
measured in liters.
96
Procedure
Before coming to lab, determine the molar masses of methanol (CH3OH), acetone
(CH3COCH3), and hexane (C6H14) and list them in your notebook. At the start of lab
period, record the current barometric pressure reading provided by your instructor.
Label a clean 50 mL beaker with your initials. Give it to your instructor to get an
unknown sample of one of the three volatile liquids listed above. When not transferring
the sample, cover the beaker with a watch glass. Your instructor will also loan you two
clean 125 mL Erlenmeyer flasks for this experiment -please return these to your instructor
when you have finished.
Determine the volume of each flask by filling to the top with water and measuring the
volume of the water with a graduated cylinder. Note the volume of each flask will be
slightly different, and will NOT be equal to the nominal volume of 125 ml.
After you have determined the volume of the flasks, dry them carefully with a paper
towel then dry them completely on your hotplate. If any water remains in your flasks the
last part of the experiment will not work properly.
Put approximately 400 mL of distilled water into each of two 600 mL beakers and place
these on a hotplate to boil (setting~5). Add about 2 mL of 1 M HCl (aq) into each beaker
to prevent scale formation.
While the water is heating, obtain four 3 x 3 inch squares of aluminum foil and four 6inch pieces of copper wire. Place one square over the mouth of each flask and loosely
fold the edges around the rim. Use a pin to prick as small as possible a hole in the center
of the foil. Label the flasks for identification. Record to within 0.001 g the mass of each
flask together with its foil and a piece of wire. Put 5 mL of your unknown into each flask.
Cap each flask carefully with the foil; crimp the foil around the rim of the flask and
secure with the wire by wrapping the wire snugly around the neck of the flask just under
the rim. Twist the ends of the wire together while being careful not to tear the foil.
Before the water is boiling, slowly lower each flask into a beaker until it is submerged up
to the neck and secure it there using a clamp. When the water reaches its boiling point,
reduce the hotplate setting to maintain a slow boil until no more of your unknown is
visible in the flasks. Then boil five more minutes. If the water level drops during this
time, add more to maintain the level near the top of the flasks.
Remove the flasks from their water baths, wipe them dry, and allow them to cool to room
temperature. Gently wipe off any water that remains on the foil. Weigh the flasks and
their contents to within 0.001 g.
97
Repeat the experiment, again using both flasks, so that you have a total of four trials.
Calculations and Report
In the table below, look up the boiling point temperature of water at the current
barometric pressure.
The initial mass of each flask assembly includes the mass of the air inside the flask. The
mass of this air cannot be ignored in comparison to the mass being determined for the
unknown vapor. Air is comprised of 78% N2, 21% 02, and 1 % Ar. As a result, air has an
average molar mass of M = 29.0 g/mol. By rearranging equation (2), we may calculate
the mass, m, of air in a known volume:
(6)
Given the volume of your flasks, calculate the mass of air present in each initial
weighing, and subtract this to determine the empty (no air) mass of the flask assembly for
each trial.
For each trial, subtract the mass of the empty flask from the final mass of the flask with
the unknown vapor. Use equation (2) to calculate the molar mass of the unknown for each
of your trials and average the results. Compute the standard deviation of your data set this is a measure of the uncertainty in your result.
Complete the report sheet. In view of your calculated molar mass and its uncertainty,
predict the identity of your sample. Is the standard deviation of your trials larger than the
difference between your measured molar mass and the actual molar mass of the
compound you predict? If so, how certain are your results?
98
THE VAPOR PRESSURE OF WATER NEAR IT’S BOILING POINT
p(atm)
T(Co)
p(atm)
T(Co)
p(atm)
T(Co)
0.692
0.6946
0.6973
0.6999
0.7026
0.7052
0.7079
0.7106
0.7133
0.716
0.7187
0.7214
0.7241
0.7269
0.7296
0.7324
0.7351
0.7379
0.7407
0.7435
0.7463
0.7491
0.7519
0.7547
0.7575
0.7604
0.7632
0.7661
0.7689
0.7718
0.7747
0.7776
0.7805
0.7834
0.7863
0.7892
0.7922
0.7951
0.7981
0.801
0.804
0.807
0.81
0.813
90.0
90.1
90.2
90.3
90.4
90.5
90.6
90.7
90.8
90.9
91.0
91.1
91.2
91.3
91.4
91.5
91.6
91.7
91.8
91.9
92.0
92.1
92.2
92.3
92.4
92.5
92.6
92.7
92.8
92.9
93.0
93.1
93.2
93.3
93.4
93.5
93.6
93.7
93.8
93.9
94.0
94.1
94.2
94.3
0.816
0.819
0.822
0.8251
0.8281
0.8312
0.8342
0.8373
0.8404
0.8435
0.8466
0.8497
0.8528
0.856
0.8591
0.8623
0.8654
0.8686
0.8718
0.875
0.8781
0.8814
0.8846
0.8878
0.891
0.8943
0.8976
0.9008
0.9041
0.9074
0.9107
0.914
0.9173
0.9206
0.924
0.9273
0.9307
0.934
0.9374
0.9408
0.9442
0.9476
0.951
0.9545
94.4
94.5
94.6
94.7
94.8
94.9
95.0
95.1
95.2
95.3
95.4
95.5
95.6
95.7
95.8
95.9
96.0
96.1
96.2
96.3
96.4
96.5
96.6
96.7
96.8
96.9
97.0
97.1
97.2
97.3
97.4
97.5
97.6
97.7
97.8
97.9
98.0
98.1
98.2
98.3
98.4
98.5
98.6
98.7
0.9579
0.9613
0.9648
0.9683
0.9718
0.9753
0.9788
0.9823
0.9858
0.9893
0.9929
0.9964
1
1.0036
1.0071
1.0107
1.0143
1.0179
1.0216
1.0252
1.0288
1.0325
1.0362
98.8
98.9
99.0
99.1
99.2
99.3
99.4
99.5
99.6
99.7
99.8
99.9
100.0
100.1
100.2
100.3
100.4
100.5
100.6
100.7
100.8
100.9
101.0
99
Name:
Report Sheet
Molar Mass of a Vapor
gas constant, R (include units)
barometric pressure, P (from instructor)
boiling temperature, T (from table)
Unknown (A, B or C) _______________________
DATA
(Sample data)
trial 1
trial 2
flask volume (mL)
mass of dry flask (including
air) (g)
mass of air in flask (g)
mass of dry flask (no air) (g)
mass of flask + vapor (g)
mass of vapor (g)
molar mass of vapor (g/mol)
Show how you calculated:
The mass of air for trial 1
The molar mass of the unknown vapor for trial 1
100
trial 3
trial 4
Average molar mass =
Standard deviation =
Identity of unknown sample:
101
Experiment 14.
Thermochemistry
Purpose
i
i
i
In this experiment, you will:
use a calorimeter to measure the temperature change which
occurs during a chemical reaction
calculate the heat energy evolved by the reaction compare
reaction enthalpies for four acid-base reactions
use experimental results to calculate the formation enthalpy
of the acetate ion
Background
When a chemical reaction occurs in a system at constant pressure under conditions such
that no energy is lost to or gained from the surroundings, it is generally found that the
temperature of the system either increases or decreases. When the system is in thermal
contact with the surroundings, energy will be lost as heat if the temperature rises due to
reaction and the reaction is said to be exothermic. If there is a decrease in temperature of
the system due to reaction, energy will be gained from the surroundings as heat and the
reaction is said to be endothermic.
Heat changes that take place in a system at constant pressure are related to a property of
the system called enthalpy, denoted by the symbol H. The SI unit of enthalpy is the same
as that of energy, joule (J) Enthalpy is a state function, and can be related to heat changes
at constant pressure by:
heat absorbed = increase in enthalpy = ÄH = Hfinal - Hinitial
(1)
When a system is configured such that no heat is gained or lost, the system is said to be
under adiabatic conditions. Thus for adiabatic processes at constant pressure, the change
in the enthalpy of the system equals the heat absorbed by the system, which is zero. Since
enthalpy can be changed by a change in temperature of the water (ÄT H2O) or by a
chemical reaction or by both, a chemical reaction carried out under adiabatic conditions
would result in the following relationships:
ÄH system = ÄH ÄT H2O + ÄH due to reaction = 0 (2)
ÄH due to reaction = - ÄHÄT H2O (3)
102
The change in enthalpy due to temperature changes can be calculated by:
ÄH ÄT H2O = specific heat capacity × mass × ÄT (4)
where ÄT = Tf - Ti and can be positive (for exothermic reactions), or negative (for
endothermic reactions). If the change in temperature is measured for a chemical reaction
carried out under adiabatic conditions, the enthalpy change due to the reaction can be
determined from the enthalpy change causing the temperature change.
Enthalpy is also an extensive property, thus depending on the amount of reaction that
occurs. To obtain an intensive property for comparing reactions, we normally relate the
enthalpy change to the amount of one of the reactants or products (e.g., species A). Thus a
reaction enthalpy is given as:
(5)
The SI unit of ÄH rxn is J mol-1 , although reaction enthalpies are typically reported in kJ
mol-1 . Since we need the amount of A reacted, it is useful to select the limiting reagent
as species A.
We have seen in lecture that reaction enthalpies can be calculated from tabularized values
of standard molar enthalpies of formation, from tabularized combustion enthalpies, and
from experimental calorimetry data. It is this last method which will be used in this
experiment using the principles derived above.
The Calorimeter
A calorimeter is a device with which we can measure the enthalpy change in a system
during a reaction by monitoring the temperature increase (for an exothermic reaction) or
decrease (for an endothermic reaction). Several types of calorimeters have been
demonstrated in the text and in lecture. In this experiment, the calorimeter used is a
Styrofoam capped cup. The foam between the liquid and surroundings acts as a thermal
barrier to prevent heat loss from the reaction to the surroundings. The reactions will be
carried out in aqueous solution in the cup, and the temperature will be monitored with a
standard laboratory thermometer. We will make two major assumptions about the
calorimeter and system. First, the device will be considered an ideal calorimeter,
operating adiabatically at constant pressure; i.e., all of the heat evolved by the reaction
goes to raising the temperature of the reaction solution, and none is lost to the beaker and
air space. Second, we assume that the final solution is sufficiently dilute so that its density
and specific heat capacity will be those of pure water (1.00 g mL-1 and 4.184 J g–1 C-1,
respectively). Both of these assumptions will introduce some error, and one report
103
question will require the calculation of this error.
The enthalpy change associated with the temperature change is calculated from equation
(4) above. From this result, the enthalpy change due to the reaction can be obtained from
equation (3). This result can be converted to kJ and then, after the amount of limiting
reagent is determined, the AHrxn can be determined from equation (5).
Chemistry of the Reactions
The reaction enthalpy for four acid-base reactions will be determined:
I.
II.
III.
IV.
HCl (aq) + NaOH (aq) 6 NaCl (aq) + H2O (1)
HCl (aq) + NH3 (aq) 6 NH4Cl (aq)
HC2H3O2 (aq) + NaOH (aq) 6 NaC2H3O2 (aq) + H2O (l)
HC2H3O2 (aq) + NH3 (aq) 6NH4C2H3O2 (aq)
HCl and NaOH are strong (acid and base, respectively), while HC2H3O2 and NH3 are
weak. In aqueous solution, the weak systems do not to ionize significantly, so the net
ionic equations for the above reactions reduce to:
I.
II.
III.
IV.
H3O+ (aq) + OH- (aq) 6 2H2O (1)
H3O+ (aq) + NH3 (aq) 6 NH4+(aq) + H2O(l)
HC2H3O2 (aq) + OH- (aq) 6 H2O (1) + C2H3O2- (aq)
HC2H3O2 (aq) + NH3 (aq) 6NH4+ (aq) + C2H3O2- (aq)
As Bronsted-Lowry acid-base reactions, all involve a proton transfer from the add to the
base. More stable bonds are formed in these reactions, so the reactions are, therefore,
exothermic.
104
Procedure
Arrange your styrofoam cups, thermometer and cap as shown by the instructor. Clean,
dry, and label two 50 mL and two 100 mL graduated cylinders (one from each partner)
for the four reactant solutions to be used in this experiment (this will prevent crossmixing of solutions).
Use the following procedure to measure ÄT:
1. Obtain 45.0 mL each of the acid and base needed for the run in the designated
graduated cylinders.
2. Add the acid to the styrofoam cup calorimeter, and then measure and record the initial
temperature (Tinitial ).
3. Rapidly pour the base into the cup, replace the cover, and stir the reaction solution
with calorimeter stirrer.
4. Observe the temperature of the reaction mixture while continuously stirring, and
record the highest reading as the final temperature (Tfinal ).
5. When finished with a measurement, pour out the solution, rinse and dry the cup,
stirrer, and thermometer, and perform the next run.
Experiment I [Run three times]
45.0 mL 2.00 M HCl (aq) + 45.0 mL 2.00 M NaOH (aq)
Experiment II [Run three times]
45.0 mL 2.00 M HCl (aq) + 45.0 mL 2.00 M NH3 (aq)
Experiment III [Run three times]
45.0 mL 2.00 M HC2H3O2 (aq) + 45.0 mL 2.00 M NaOH (aq)
Experiment IV [Run three times]
45.0 mL 2.00 M HC2H3O2 (aq) + 45.0 mL 2.00 M NH3 (aq)
105
Calculations and Report
In the following calculations, assume that the specific heat capacity for the reaction
mixture is 4.1841 J g-1 C-1 and its density is 1.00 g mL-1. For each run, calculate ÄT,
change in enthalpy due to the reaction (in kJ), amount of limiting reagent, and the
reaction enthalpy (in kJ/mol limiting reagent). Show the calculation setup for run 1 of
Experiment I as indicated on the report sheet. Calculate the average reaction enthalpy for
each of the four experiments and then answer the questions.
Questions for your Notebook
1. How many grams of HCl are in 45 mLs of 2M HCl?
2. How many grams of NaOH are in 45 mLs of 2M NaOH?
3. Why can you assume that your solution contains 90 mLs of water with a heat capacity
of 4.184 J g-1 C-1 and ignore the other ions in the solution?
4. Also make a diagram of a typical experimental set up.
106
Name: _________________
_________________
Report Sheet
Thermochemistry
Experiment I (HCl/NaOH)
Run 1
Run 2
Run 3
Initial Temperature, Ti(oC)
_______
______
_______
Final Temperature, Tf(oC)
_______
_______
_______
_______
_______
_______
ÄHH2O (kJ)
_______
_______
_______
ÄH due to rxn (kJ)
_______
_______
_______
_______
_______
_______
_______
_______
_______
ÄT(oC)
Amount limiting reagent (mol)
ÄHrxn (kJ/mol)
Average ÄHrxn __________________kJ/mol
In the space below show the sequence of calculations for run 1:
107
Run 1
Run 2
Run 3
Initial Temperature, Ti(oC)
_______
_______
_______
Final Temperature, Tf(oC)
_______
_______
_______
ÄT(oC)
_______
_______
_______
ÄHH2O (kJ)
_______
_______
_______
ÄH due to rxn (kJ)
_______
_______
_______
_______
_______
_______
_______
_______
_______
Experiment II (HCl/NH3 )
Amount limiting reagent (mol)
ÄHrxn (kJ/mol)
Average ÄHrxn __________________kJ/mol
Run 1
Run 2
Run 3
Initial Temperature, Ti(oC)
_______
_______
_______
Final Temperature, Tf(oC)
_______
_______
_______
ÄT(oC)
_______
_______
_______
ÄHH2O (kJ)
_______
_______
_______
ÄH dur to rxn (kJ)
_______
_______
_______
_______
_______
_______
_______
_______
_______
Experiment III (HC2H3O2 /NaOH)
Amount limiting reagent (mol)
ÄHrxn (kJ/mol)
Average ÄHrxn __________________kJ/mol
108
Run 1
Run 2
Run 3
Initial Temperature, Ti(oC)
_______
_______
_______
Final Temperature, Tf(oC)
_______
_______
_______
ÄT(oC)
_______
_______
_______
ÄHH2O (kJ)
_______
_______
_______
ÄH due to rxn (kJ)
_______
_______
_______
_______
_______
_______
_______
_______
_______
Experiment IV (HC2H3O2 /NH3 )
Amount limiting reagent (mol)
ÄHrxn (kJ/mol)
Average ÄHrxn __________________kJ/mol
Answer the following questions
1. Would I get a different ÄHrxn in experiment I if I used 45.0 mL of 2.00 M HNO3
instead of HCl? Explain why or why not.
109
2. Given the following data, calculate the theoretical values for ÄHrxn for experiments I
and II. Compare these values with your experimental results for ÄHrxn by calculating the
percentage error for each experiment (See Appendix 1.)
species
ÄHf(kJ/mol)
H3O+(aq)
-285.83
H2O(l)
-285.83
OH-(aq)
-229.99
NH3(aq)
-80.29
NH4+(aq)
-132.51
HC2H3O2
-485.76
3. Using your average result for experiment III, calculate ÄHfo for the acetate ion. Show
your setup.
110
Experiment 15.
Determination of Glucose using a Spectrophotometer
Purpose: In this experiment you will:
i
Oxidize glucose to gluconic acid, while reducing Fe(CN)63- to Fe(CN)64-.
i
Prepare a standard curve showing how the absorbance of Fe(CN)63decreases with increase glucose concentration.
i
Determine the concentration of an unknown glucose sample using this
standard curve.
Background
Glucose, also known as Dextrose, is a simple sugar with the empirical formula C6H12O6.
It forms important biopolymers such as starch, cellulose, and glycogen. Carbohydrates
are absorbed in the blood stream as glucose and the sugar is oxidized by the body to
produce energy. The body has a hormonal system that tries to regulate the level of
glucose in the blood. Normal levels of blood glucose are between 50 and 140 mg/dl (2.7
- 8 mM). If your blood sugar drops much below this, you can go into hypoglycemic
shock, which, in turn, can lead to death. On the other hand, if your blood sugar is
consistently above the 140 mg/dl level, you may have diabetes, and there are a host of
problems associated with this state as well.
Pharmacies sell over-the-counter kits designed to help diabetics monitor their blood sugar
level. These are simple colorimetric assays in which you put a drop of blood on a
specially treated plastic strip and watch a color change occur due to different chemical
reaction that occur on the test strip. These test strips are somewhat expensive and not
terribly accurate. In a clinical setting, like a hospital, glucose testing is done routinely
and more accurately using a variety of different chemical reactions. The lab we do this
week modeled after one of these reactions.
The method used here is based on the oxidation of glucose by ferricyanide ion,
Fe(CN)6-3. The net reaction is:
H2O + C6H12O6 + 2 Fe(CN)6-3
Glucose
Ferricyanide
Yellow
6 C6H12O7 + 2 H+ + 2 Fe(CN)6-4
Gluconic Acid
Ferrocyanide
Colorless
111
Ferricyanide ion has an absorbance maximum at 420 nm and so will have a maximum
absorbance at that wavelength. However, if a sample contains both ferricyanide and
glucose, the above chemical reaction will occur, and the amount of ferricyanide in the
solution will be reduced, so the absorbance of the solution will go down.
In this experiment a series of solutions are made that contain a constant amount of
ferricyanide, but increasing, known, amounts of glucose. The linear relationship between
absorption at 420 and glucose concentration can be used to make a Standard Curve that
correlates the absorbance of the solution to with the amount of glucose in the solution.
An unknown solution of glucose will then be analyzed using the same concentration of
ferricyanide and the same chemical reaction. The absorbance of this solution can be
plotted on the standard curve, and the original concentration of glucose in the unknown
determined from the standard curve.
Procedure
You will find the following reagents and materials in the lab:
(A) 0.015 M Fe(CN)6-3, 0.5M Na2CO3 (SOLUTION A)
(B) 5.000 mg/100ml Glucose standard (SOLUTION B)
(C) 50 mL or 100 mL volumetric flasks
(D) unknowns
Preparation of standard curve (1 curve per group)
Introduce 4.00 mLs of solution A into each of four clean 100 mL volumetric flasks. Mark
the flasks 1, 2, 3, and 4. Now pipet 10.00 mL of solution B into flask 2, 20.00 mL of
solution B into flask 3, and 30.00 mL of solution B into flask 4.
Half fill each of the flasks with water. Place in a boiling water bath for 15 minutes.
Make sure they are unstoppered. (also make sure the flasks don't tip). Cool the flasks
to room temp and fill to the mark with distilled H2O. Cap and thoroughly mix the
contents of the flasks.
Using distilled water as the reference, measure the absorbance of each of the standard
solutions and the unknown at 420 nm. For the set of standards, plot the absorbance vs.
the glucose concentration.
112
Preparation of the unknown (1 unknown for each person)
Introduce 2.00 mL of solution A into a 50 mL volumetric flask. Pipet 20.00 mL of your
unknown into the flask. Half fill the flask with water. Place in a boiling water bath for
15 minutes. Make sure it is unstoppered. (Also make sure the flask doesn't tip). Cool
the flask to room temp and fill to the mark with distilled H2O. Cap and thoroughly mix
the contents of the flasks. Determine the absorbance of the flask at 420 nm, determine the
concentration of the original Glucose unknown using your standard curve.
113
Name:
Report Sheet
Determination of a Glucose Unknown
I. The standard curve
A. Concentration of glucose in Solution B ___________(mg/100mL)
B. Concentration of glucose in
Flask 1___________(mg/100ml)
Flask 2___________(mg/100mL)
Flask 3___________(mg/100mL)
Flask 4___________(mg/100mL)
C. Calculation used to determine glucose concentration in flask 3
D. At 420 nm
Flask 1
Transmittance
__________
Absorbance
__________
Flask 2
__________
__________
Flask 3
__________
__________
Flask 4
__________
__________
114
E. Standard Curve
II. Determination of Unknown
A.
Name
of student
Unknown
Number
Absorbance
of Unknown
Concentration of Unknown (mg/100mL)
Dilute solution
Original Solution
________
______
_________
____________
______________
________
______
_________
____________
______________
________
______
_________
____________
______________
________
______
_________
____________
______________
B. Sample calculation for converting from concentration of unknown in dilute solution to
concentration of unknown in original concentrated solution.
115
Experiment 16.
Physical Properties of Chemicals - Melting Points, Boiling
Points and Sublimation
Purpose: In this experiment you will:
i
Determine the melting point of chemicals
i
Determine the boiling point of chemicals
i
Demonstrate the sublimation of a solid
Background
Just as every object in a room can be described in terms of its physical properties, every
chemical has can be characterized by a number of different physical properties. While a
given physical property may not uniquely identify a chemical, a complete set of physical
properties can often be used to identify a compound unambiguously. Many different
handbooks are available to chemists that catalog the known physical properties of
chemicals. These handbooks not only help a chemist know what will happen to a given
chemical as he manipulates it, but also lets the chemist identify unknown compounds.
Some of the most common physical properties used to characterize chemicals include
physical form (solid-liquid-gas), color, smell, solid form (metal-crystal-amorphous),
density, refractive index, melting point, boiling point, and solubility in various solvents.
These properties not only help to identify a compound, but deviations from ideal values
can be used to assess purity. Whenever a new chemical is isolated in the lab, a large
amount of time is spent determining any or all of the above properties to help to
characterize the new compound. A complete characterization of the compound is a
required part of reporting this compound in the literature.
One of the most common physical properties reported for a solid is its melting or freezing
point. Our text refers to this as the normal melting point, or the temperature at which a
solid and liquid state of a compound have the same vapor pressure under conditions
where the total atmospheric pressure is 1 atm. While this definition sound pretty
formidable, what it really means is the temperature at which the solid and liquid coexist
(a slush forms) at normal atmospheric pressure.
The melting points of solids depend on the nature of the solid. Ionic solids generally have
melting points in the several 100's to 1000's oC, while organic compounds held together
116
with a mixture of London dispersion forces and dipole-dipole forces will melt somewhere
near room temperature (±100o C). Finally small gases that are held together only by
London dispersion forces have melting temperature in the -100o to -250o C range.
A typical melting temperature for a pure compound has a very small range, it can be as
little as 1/10 of a oC in range. In the general lab setting, a melting range of 1o C is typical
and indicative of a relatively pure substance. Melting temperature ranges larger than this
indicate that the substance is not pure.
Similarly the normal boiling point, is often used to characterize liquids. The definition
of the normal boiling point is the temperature at which a liquid has a vapor pressure of
exactly 1 atm. Since a liquid boils when its vapor pressure equals the atmospheric
pressure, this definition usually mean simply the temperature at which the liquid boils.
Doing boiling point determinations in Spearfish presents one small problem. We are
working at a high enough altitude that our air pressure is seldom 1 atm! Typically our air
pressure is on the order of .9 atm, so water, which is supposed to boil at 100oC will boil
here at 96oC. The bottom line here is not to be too surprised if all your boiling point
determinations today turn out to be a bit low.
Sublimation refers to the direct conversion of a solid to a gas, without going through in
intermediate liquid state. It can also refer to the reverse process, a gas converting directly
to a solid without going through the liquid state. Relatively few compounds can do this;
familiar examples include naphthalene, caffeine, iodine, solid CO2, and water.
Sublimation does not have as distinctly defined a temperature like the melting or boiling
points, but it is still a useful physical property, both to identify a compound, and
sometimes to purify it.
Procedure
Experiment 1
Determination of a melting point
Obtain about 0.1 g of an unknown compound from the instructor. (Precise weight is not
required, just eyeball an amount that looks about right)
Also obtain a capillary tube. If it has not already been done, cut the capillary tube in half
with a file.
What you are trying to do is to get crystals of your unknown into the bottom closed
end of the capillary tube to a depth of about 3 mm. If the crystals of your unknown are
large you may need to obtain a mortar and pestle to grind the crystals down to a smaller
size. Otherwise simply force some of the unknown into the open end of the tube by
117
gently pushing the tube straight up and down into the unknown, then turn the tube endfor-end and tap it on the table top to get the crystals to fall to the bottom of the tube. If
you need it, you can also drop your capillary tube down a larger tube to pack the crystals
into the bottom (See demonstration).
Alternatively, if you are using the gray melting point apparatus with the thermometers
that are mounted horizontally, all you need to do is to place a sample between two cover
slips and place this ‘sandwich’ on the heating block of the machine.
The basic experiment is simply to heat the tube and report the temperature at which the
solid melts. This is usually reported as a range, with the lower temperature corresponding
to the point where the crystals of the solid first start to melt and have a ‘slushy’
appearance, and the upper temperature corresponds to the temperature where the last of
the solid completely disappears. As you might expect, you get a smaller range and
sharper melting point if you heat the tube slowly. On the other hand you can spend all
day waiting for things to happen if you don’t heat fast enough. You’ve got to find the
happy medium rate of heating. Another way this can be done is to prepare two samples.
Heat the first sample quickly to get the temperature of the bath in the right range, then
cool the bath slightly, put your second sample in and heat the material slowly through its
melting range.
When finished the unused chemicals may be disposed of in the garbage cans, while the
capillary tubes should be dropped in the broken glass can.
118
POSSIBLE UNKNOWNS and their melting points (oC)
Acetamide
Acetanilide
Benzophenone
Benzoic Acid
Biphenyl
Lauric Acid
Naphthalene
Stearic Acid
82
114
48
121
70
43
80
70
Experiment 2
Sublimation
In the hood you will see a simple apparatus set up to demonstrate sublimation. It consists
of solid naphthalene placed in the bottom of a 100 ml beaker, and then a second 50 ml
beaker containing ice is carefully placed inside the larger beaker so the bottom of the icecold beaker is suspended just above the naphthalene. The beakers are then placed on a
hot plate, and the plate is warmed just enough to encourage the naphthalene to go into the
vapor phase, but not enough to actually melt the material. As the naphthalene vapors hit
the bottom of the cold beaker they come out of the vapor phase and condense to form
solid crystals, that are usually very large, flat, and beautiful.
Let the apparatus cool, then remove the beaker as scrape some of the sublimed
naphthalene off the beaker for closer examination. Obtain some of the crude and some of
the sublimed naphthalene for melting point analysis. Was there any difference in the
range of melting points for these two materials? Note: It is usually very difficult to get
the sublimed crystals into a capillary tube, so it is best to test these crystals using the
apparatus that squashes the crystals between cover slips.
Make a diagram in you lab notebook of the apparatus used to perform sublimation.
119
Experiment 3
Determination of the boiling point of a liquid
CAUTION ALL ORGANIC SOLVENTS USED IN THE LAB ARE FLAMMABLE.
ALL BUNSEN BURNERS MUST BE EXTINGUISHED BEFORE BEGINNING
THIS PORTION OF THE LAB. See instructor for alternate places for setting up this
lab if you are ready early.
-Obtain about 5 mLs of liquid unknown from the instructor.
-Place this sample in a large test tube and place a boiling chip in the bottom of the test
tube.
-Place a 250 mL beaker full of water with a boiling chip or two on your hot plate. Set up
an apparatus stand to hold your test tube of unknown upright in the beaker and hold your
thermometer upright above the space above the liquid in your test tube. A rubber stopper
with a split may also be used to hold the test tube above the liquid
- Heat the water and observe the temperature in the air space above the liquid. The
temperature should rise steadily, and then stabilize at the boiling temperature of your
liquid. Record your boiling temperature and determine the identity of your unknown.
- Make a diagram in your lab notebook of the boiling point apparatus.
UNUSED UNKNOWNS SHOULD NOT BE POURED DOWN THE SINK.
Instead deposit in ORGANIC WASTE VESSEL.
120
Possible Liquids unknowns and their Boiling points (oC at 1atm)
Acetone (Propanone)
56
Cyclohexane
81
Ethyl Acetate
77
Hexane
69
Isopropyl alcohol (2-Propanol)
83
Methyl alcohol (Methanol)
65
1-Propanol
97
Experiment 4
Determination of Infrared (IR) and Nuclear Magnetic Resonance (NMR) absorption
spectra
While physical measurements like boiling points and melting points can be used to
characterize a compound, they are not very good at identifying unknown materials. There
are two separate problems. One is that may compounds have the same, or nearly the
melting point or boiling point. The other is that even the modest altitude that we have in
Spearfish is enough to through the boiling point off by a degree or two.
Two other techniques that are used routinely by organic chemists to characterize and
identify unknown materials are called Infrared (IR) and Nuclear Magnetic Resonance
(NMR) spectroscopy. While the detailed explanations of how these techniques work is
beyond the scope of this class, a brief introduction to both of these techniques is included
as supplemental material and can be found in this lab manual immediately following the
report sheet. Read both of these supplements.
121
Near the IR machine you will find bottles labeled A-G. Find the bottle that matches the
unknown you used in the boiling point experiment. Have the instructor or TA who is
working with the IR show you how to obtain an IR spectrum for your sample. Include a
copy of this spectrum with your lab write-up and your lab noteboook. By matching this
spectrum with the spectra found in the supplemental materials you should be able to
identify your boiling point unknown.
Near the NMR machine you will find 7 pre-made samples tubes with labels A-G. Find
the tube that matches the unknown you used in the boiling point experiment. Have the
instructor or TA who is working with the NMR show you how to obtain an NMR
spectrum for your sample. Include a copy of this spectrum with your lab write-up and
your lab notebook. By matching this spectrum with the spectra found in the supplemental
materials you should be able to confirm the identity of your unknown.
122
Name:
Report Sheet
Physical Properties of Solids
Experiment I
1. Solid Unknown number___________
Melting Begins
Run I
Run II
o
_________ C________oC
Melting Ends
_________oC________oC
Melting Range
_________oC________oC
Identification of unknown _____________________
Experiment II
_________________oC
Melting range of crude naphthalene
Melting range of sublimed naphthalene _________________oC
Experiment III
Liquid unknown number
_______________________
Run I
Observed Boiling point
___________oC
Run II
__________oC
Identity of unknown: ______________________________
Structure of unknown:
Experiment IV
Attach labeled copies of all spectra obtained for your unknown.
123
Supplemental Material - Nuclear Magnetic Resonance
(NMR)
Nuclear Magnetic Resonance (NMR) is a way of characterizing an atom based on how its
nucleus interacts in a magnetic field. It is used in organic chemistry to as a way to identify
known organic chemicals and to determine the structure on novel compounds. It is used in
medicine as the basic interaction that a magnetic resonance imaging (MRI) machine uses to
obtain three dimensional images of soft tissue.
Certain nuclei have the property that they interact with an external magnetic field as if the nuclei
themselves are small magnets. Fortunately 1H (a proton) is one of those nuclei, and it is found in
virtually every organic compound. When this nucleus acts like a magnet, it lines up with an
external magnetic field just like any magnet would. The nucleus stays in alignment until it is hit
with microwave radiation of a very specific frequency. When the nucleus sees microwave
radiation with this specific frequency, the sample absorbs this radiation, and the energy that it
absorbs is used to knock the nucleus out of alignment. With some clever design we can make an
instrument that holds our sample in a magnetic field, irradiates the sample with microwave
radiation, and measures what specific frequencies of the radiation are absorbed by the sample.
All the NMR instrument is, is a very big magnet, some mechanisms to hang your sample tube
right in the very strongest part of the magnet’s field, some electronics to tickle your sample with
a touch of microwave radiation, and some additional electronics to see how much and what
frequency of radiation the sample absorbs when this energy hits it.
Like the other absorption techniques you have already seen, (Light and Infrared absorbance),
NMR data is displayed as plot of absorbance in the Y-axis versus a measure of wavelength or
frequency on the X-axis. Thus each peak in your spectrum represents a different frequency
where your sample absorbs radiation.
The magnet in our machine makes protons absorb radiation with a frequency of 90,000,000
Hertz or 90 MHz. If that was all that there was, this technique would not be very useful because
all protons would absorb radiation of the same frequency and there would be no way to tell them
apart. What actually happens in a compound, is that the electrons around each atom interact
with the magnetic field to slightly shield the nucleus from the magnetic field. Depending where
a proton is in a compound, this shielding can be larger or smaller, with protons on CH3 groups
being very shielded and protons in aromatic systems being very de-shielded. The shielding
effect is actually very, very tiny, and the frequency of the absorbed radiation is only 1 part in
1,000,000 or 1 part per million (ppm) different from our original frequency. Thus, if we have a
proton absorbing energy at 90,000,000 Hz, and we see a typical 1 ppm shielding, its resonance
frequency will vary by 1 part in a million or 1.000001, so it will resonate at 90,000,090 Hz.
These frequencies are very cumbersome to write down. To make things even more complicated,
lots of different magnets built into NMR machines, and these magnets can make protons absorb
radiation anywhere from 60 MHz up to 1,000 MHz. So we can directly compare the data
124
obtained from these machines, chemists have adopted the convention of simply setting one
resonance to be a 0 reference, and then displaying the other frequencies in ppm (parts per
million) difference from this standard. Thus in our NMR spectrum, the X–axis is measured in
ppm. Further, when we prepare your sample, we use a solvent with the reference standard
already in the solvent, so you will have a peak in your spectrum at 0 ppm, and you will use this
peak to calibrate you X-axis. Once this calibration is done in the computer, all the other peaks
are set, any you can directly compare your data with that taken on any other instrument,
regardless of how big or small it’s magnet was. Since our reference contains extremely well
shielded protons, we place this zero on the right hand side of our spectrum. Our sample usually
contains protons that are less shielded so you see the protons of your sample as peaks to the left
of the reference, and we use a chemical shift scale running form right to left to display how
many ppm difference there is between your sample peaks and the standard.
During the day’s experiment you might notice another reason for using this internal standard..
Our machine uses a permanent magnet, and the magnetic field actually changes and drifts around
slightly over the course of the day. We used a standard sample in the morning and set the
machine so the standard was at zero initially, but, by the time you run your sample, there is a
good chance that the field has drifted slightly so your reference peak is no longer at 0.00 ppm. If
you notice this, make sure you are shown how to readjust the data display so your reference is
back at zero where it should be so all your other peaks come out where they are supposed to.
While the interactions that shield and deshield a nucleus can be calculated, these calculations are
way beyond what most chemists want to deal with, so we will use our X axis in a very empirical
manner. We know that under normal circumstances most protons in a sample resonate between
0 and 10 ppm. Protons in CH3 groups are usually around 1 ppm, protons in CH3 groups next to a
C=O are a but further to the left (have a higher ppm number), and protons in an aromatic ring
system will have ppm values in the 6-10 region.
You will see in your NMR spectrum that there is also lots of fine structure, or smaller subpeaks
within each main peak. To a trained organic chemist these peaks all have special meaning and
can be used to completely identify every single resonance in a spectrum and to associate it with
distinct protons in a chemical structure. We won’t go that far in this lab. Here we will just
remember that every compound has a unique NMR spectrum, and that this spectrum is a
constant and does not change, so if you look at a pattern from a pure compound, and then see
that same pattern in your sample, you know that you have that compound in your sample.
Below are NMR spectra for the solvents hexane, methanol, isopropyl alcohol, ethyl acetate,
cyclohexane, acetone, and 1-propanol, the unknowns used in your boiling point experiment.
Near the NMR you will also find a folder containing these spectra presented in a larger format,
and with a diagram relating peaks in the NMR spectrum to protons on the molecule’s structure.
125
Using boiling points to identify compounds is not the greatest way to determine the identity of a
material. Lots of different compounds will have the same boiling point, and impurities can make
the boiling point change. On top of that, in Spearfish our elevation makes liquids boil 2-3
degrees lower than expected, so you will probably identify your unknown incorrectly.
This is where NMR comes in. The NMR spectrum of every compound is unique, and it can be
used to absolutely identify an organic compound. At the NMR you will find 7 sample tubes
already filled with the solvents you used in the boiling point experiment. Find the tube that
matches the unknown you used. Place it in the NMR instrument and obtain an NMR spectrum.
Once you have printed out the spectrum, compare it to the reference spectra and identify your
liquid unknown. Was your guess based on the boiling point correct?
Reference Spectra
1. Acetone
126
2. Cyclohexane
3. Ethyl Acetate
127
4. Hexanes
5. Isopropyl Alcohol
128
6. Methyl alcohol
7. 1-Propanol
129
Supplemental Material - Infrared (IR) Spectroscopy
Infrared (IR) Spectroscopy is another method used to characterize compounds in the chemistry
lab. It is exactly analogous to the visible spectroscopy we did in Experiments 10 and 11, the
only difference is that we are using low energy infrared light instead of higher energy, visible
light.
When you excite molecules with IR energy, the added energy is not enough to push electrons
from one orbital into another, like we were doing with visible light. Instead what you are doing
is making the atoms in the molecule vibrate back and forth in different ways. Since every
compound has a different set of chemical bonds to rotate and vibrate, the IR spectrum of every
compound is unique like a fingerprint and can be used to absolutely identify a compound.
Just as we did in Experiment 10, the first step is to scan a compound’s absorbance at many
wavelengths, and then to plot this information as intensity vs wavelength. At typical IR
spectrum is shown below:
Hexanes (straight and branched chains)
You should notice two main differences between this curve, and the one you obtained for the dye
in experiment 10.
First, the Y axis is in % transmission instead of absorption. IR techniques are not nearly as
quantitative as UV techniques, hence most of the time you don’t bother with converting your
signal from raw % transmittance to true absorbance
Second, the X axis is in units of wavenumber (cm-1) instead of wavelength. If you go back to
Experiment 9, you will see that wavenumber is a measure of frequency, and when wavenumber
is in units of cm-1 this number means how may waves you would find in one centimeter.
Historically these were the first units used in IR spectroscopy, and they have remained the
preferred unit. You should remember, however, that it is just a little algebra to convert from
frequency to wavelength so the units are, essentially interchangeable.
130
Now that you have seen the differences, let’s talk about the spectrum itself. First, remember that
a high transmittance means that the light is being transmitted though the sample and it is not
being absorbed. Thus, in this kind of plot, you should have mostly signal along the upper edge,
indicating that light is passing through the sample, and an absorbance corresponds to a valley or
trough that goes down. In the hexane sample the only chemical groups we have are -CH3 and CH2-. You can see one major valley around 2800-3000cm-1. This occurs from the C-H bond
stretching and deforming in different ways and is an indicator of the presence of CH moieties in
a compound. There are also more absorbances in the 1450-1500, and 1350-1400 range that arise
from more intricate deformations of the H-C-H bonds
Compare the straight chain hexane spectrum to that of cyclohexane below:
Cyclohexane
Notice how you still have bands at 2800-3000, 1450-1500, and 1350-1400 cm-1, but how they
shift around a little and the relative intensities have changed slightly. While cyclohexane is still
a hexane, the carbons are restrained to the ring system, so the vibrations that can occur within
this system are slightly different. Thus you get a spectrum that shows all the characteristics of a
hydrocarbon, yet has some unique differences that can be used to identify an exact compound.
One word of caution in comparing spectra, never use the absolute depth of an absorbance to
compare spectra, only use the frequencies or positions at which the absorbances occur. As you
saw in Experiment 10 absorbance depends on three things, the molar absorbtivity, the
concentration and the path length. In this experiment where we are simply putting a few drops
between two windows and squeezing the window together, the pathlength varies from
experiment to experiment, so the absolute absorbance, and the depth of the peak will also vary.
Now lets look at some other functional groups
Alcohols
Below are the spectra of three different alcohols, Methanol (CH3OH), 1-propanol
(CH3CH2CH2OH) and Isopropanol (CH3CHOHCH3)
131
Methanol
1-Propanol
Isopropanol
In addition to the CH bands we saw earlier, there are now several added peaks. Notice the large,
broad peak in the 3500 range this absorption is due to the stretching of the O-H bond. In
addition there are one or more sharp bands around 1000 cm-1. This band is due to C-O stretching
and OH deformations. It is usually around 1050 for 1o alcohols, 1100 for 2o alcohols, and 1150
for 3o alcohols.
132
A C=O bond is going to have a different vibration frequency than a C-O-H bond, so let’s
examine the spectrum of a ketone, in this case acetone.
Acetone
Notice how the broad OH band around 3500 has disappeared? The strong absorbance around
1700cm-1 is a C=O stretching band that is diagnostic for the presence of a ketone. The frequency
of this band shifts slightly when other constituents are located near the ketone. The sharp but
weaker band at 1400 is also associated with the presence of a ketone
Finally let’s look at an ester, ethyl acetate, since this compound will include both C=O stretches
and C-O-C stretches:
Ethyl Acetate
Notice how the C=O stretch has shifted slightly up to around 1750cm-1. We have also picked up
a new C-O stretch in the 1250 cm-1 region. The presence of both of these strong bands is
indicative of an ester.
133
Experiment 17.
Determination of the Enthalpy of Vaporization of H2O
Purpose: In this experiment you will:
<Determine the vapor pressure of water at several temperatures
<Use a plot of ln(Pwater) vs 1/T to determine ÄHvap for water
<Calculate the normal boiling point of water
Background
Molecules of a liquid can escape the liquid’s surface and go into the gas phase in a
process called vaporization or evaporation. To do this, energy must be supplied to the
system to overcome the forces holding the liquid molecules together. The energy
required to vaporize one mole of a liquid is called the heat of vaporization, or the
enthalpy of vaporization, ÄHvap
Given enough time, most liquids will evaporate away to nothing because some molecules
in a liquid are always escaping from a liquid’s surface. On the other hand, if you put a
liquid in a sealed container the material will not evaporate away. Why is that? In the
sealed container molecules are still escaping from the liquid’s surface, but, at the same
time, molecules are also condensing out of the gas phase and going back into the liquid.
The final result is that we reach an equilibrium state where the number of molecules
leaving the liquid is equal to the number of molecules re-entering the liquid.
If we measure the vapor pressure of the gas now, when it is at equilibrium with the liquid,
we are measuring the equilibrium vapor pressure or, more simply, the vapor pressure
of the liquid.
The vapor pressure of a liquid varies with its temperature. As the temperature of a liquid
increases, the individual molecules have more energy, so more molecules can escape
from the liquid into the vapor phase. The more molecules we have in the vapor phase, the
higher the vapor pressure. This is not a linear relationship; the vapor pressure increases
exponentially with the temperature of the liquid. The relationship between pressure and
temperature can be turned into a linear relationship if we plot the natural logarithm of the
pressure vs 1/T. Plotting the data in this manner gives us an added bonus, the slope of the
line is equal to the enthalpy of vaporization, ÄHvap. The exact relationship is:
In today’s lab we will determine the vapor pressure of water at several temperatures, plot
the ln (natural logarithm) of the pressure against 1/T and use the slope of the line to
determine the ÄHvap of water. The literature value for this parameter is 43.9 kJ/mol. Your
final task will be to use your data to determine the normal boiling point of water.
134
Experimental Procedure
A.
Record the barometric pressure, using the barometer on the wall
B.
To set up your apparatus. For today’s experiment you will need:
10ml graduate cylinder
1000ml beaker
test tube tongs
hot plate
temperature probe
You will find 1000ml beakers half filled with de-ionized water already heating on hot
plates at your bench. Fill your 10ml graduate cylinder with enough de-ionized water to
give a bubble of 3ml-4ml volume when you quickly invert the cylinder inside the 1000ml
beaker, trapping the bubble in the cylinder. Use the test tube tongs to hold the cylinder in
place in the beaker. Fill the beaker with de-ionized water, covering the graduate cylinder
completely. Heat the water to a temperature of 75o to 80oC. The bubble inside the
graduate cylinder will expand as the temperature increases, allow it to expand beyond the
gradations on the cylinder, but not to escape the confines of the cylinder. Once the water
has reached temperature carefully remove the beaker from the hot plate. Watch the
volume of the bubble carefully as the temperature drops. Once it is within the calibration
markings on the cylinder begin to record the volume of the bubble and the temperature.
Record the volume every 3oC until the water cools to 50oC.
Now add ice to the beaker to cool the water down below 5oC and record the temperature
and volume of the bubble.
Make a diagram of you experimental setup in your lab notebook.
Analysis
In looking at this experiment you should expect the volume of the gas bubble to increase
with temperature because all gases expand when they are heated. If you go through the
calculation however, you will find that the volume of the bubble gets much larger than
you would expect. This is because only a fraction of the bubble’s increase in size is due
to the expansion of the air in the bubble. Most of the expansion is due to water vapor
entering the gas bubble to further increase it’s volume.
This is where we start our calculations. At the lowest temperature, about 5oC, the vapor
pressure of water is so small that we can make the approximation that the bubble contains
only air. Since the bubble contains only air, we can go back to the ideal gas law and
calculate the number of moles of gas molecules in the bubble.
PV=nRT;
PV/RT = nair
Where nair will be the number of moles of air molecules, P will be the atmospheric
135
pressure in atm, V is the volume of the bubble in liters, R is the gas constant (.08206
L@atm/k@mol) and T is the temperature in K.
At all of the higher temperatures water will have a significant vapor pressure, and the
bubble will contain a mixture of air and water. Last semester we also learned that in
mixtures like this you can treat the two gases independently, and that the total pressure is
equal to the sums of the individual pressures. Thus:
Ptot = Pair + Pwater
We can calculate the Pair at each temperature by using the ideal gas law and solving for P
PV=nRT;
Pair =nairRT/V
We get the Pwater by recognizing that Ptot is the total pressure, and that is always equal to
the atmospheric pressure so:
Ptot = Patm = Pair + Pwater;
Pwater = Patm- Pair
At this point you now have several different Pwater values at several different temperatures
so you can start the Hvap analysis.
You have already seen the equation we are going to use:
The equation of a line is Y = mX + B, so if you plot 1/T as X, and ln(Pwater) as Y, you will
have a line with a slope = -ÄHvap/R. and the intercept of C. Please remember that the T
must be in K, and that in this equation you will use the other gas constant, R=8.3145
J/K@mol.
Plot your data as accurately as possible on your graph paper. Use a straight edge or string
to find the line of best fit. This line will have and equal number of points above it as
below it, and you will have roughly the same total distance between the points on either
side of the line and the line itself. Once you have found a line you like, draw it on your
graph. The equation for the slope of a line is (Y2-Y1)/(X2-X1), so find two convenient
points on you line, determine the X and Y values of these points, and calculate the slope
of the line. Now use this slope and the proper equation to find the Ähvap of water.
The final calculation for today’s lab is one I’m going to leave for you to figure out. The
definition of normal boiling point is the temperature at which the vapor pressure of a
liquid is exactly 1 atm. Use your data to determine this value. I can think of two
different ways to determine this using your data. Here is a hint for each method.
Method 1. Where on your plot would you find the point Pwater =1 atm?
Method 2. Look at the Clausius-Clapeyron equation in Chapter 10 of your text.
136
Name:
Report Sheet
Determination of the Enthalpy of Vaporization of H2O
Barometric Pressure ________________
V (l)
t (oC)
T (K)
1/T (K-1)
Pair (atm)
Below record volume of the bubble at the lowest temperature,
Show calculation for moles of air contained in your bubble.
Show calculation of Pwater at one temperature
137
Pwater
(atm)
ln Pwater
Graph ln Pwater vs 1/T (K), drawing the line of best fit. From your graph calculate the
slope of the line. Staple your graph to the report sheets.
Slope ____________
(You are welcome to use Microsoft excel to draw your graph, fit the best line and print
up the slope. And if you are a spreadsheet wizard you can also use the spreadsheet
formulas to calculate all of the information needed for the table on the first page of the
report sheet. Just staple a copy of your spreadsheet, showing data table and graph, to
the report sheets. )
ÄHvap for water _____________, show calculation
Normal Boiling point of water __________ , show calculation
138
Experiment 18.
Freezing-Point Depression
Purpose
In this experiment you will determine the molar mass of an unknown substance
from its freezing-point depression in a cyclohexane solution.
Background
Freezing-point depression, boiling-point elevation, and osmotic pressure are collectively
referred to as colligative properties. Colligative properties are properties that depend
only on the number of solute molecules in a solution, and not on the chemical properties
of the solute molecules.
We routinely use colligative properties of different solute in water without thinking about
it. We add antifreeze to our car’s radiator to keep the radiator fluid from freezing in
winter or boiling in summer. We throw salt on a sidewalk to lower the freezing point of
the ice on the sidewalk so it will melt.
In today’s lab we will focus on freezing-point depression. In freezing-point depression
the addition of a nonvolatile solute to a solvent to make a solution lowers the vapor
pressure of the solvent. This lowering of the solution vapor pressure both lowers the
freezing point and raises the boiling point. The exact change in the freezing point may be
calculated using the equation :
Where:
ÄT = Tfreezing point of solvent - Tfreezing point of solution.
Kf = molal freezing-point depression constant and is a characteristic of the solvent.
msolute = molal concentration of the solute (moles of solute/ kilogram of solvent).
139
Procedure
Experiment 1
Determination of the freezing point of cyclohexane
While the freezing point of a solution is easy enough to define, it is the temperature at
which both solid and liquid forms of a solute can exist, the exact determination of this
temperature can sometimes be experimentally difficult. Here we will use a cooling curve
to help us find the freezing point of our solvent. To make a cooling curve the solvent is
placed in an ice bath and the temperature of the solution recorded over time. The
temperature of the solution initially falls at a steady rate, then, as the solution freezes, the
temperature of the solution remains constant until the solvent has completely frozen.
Sometimes supercooling may be observed in the solvent actually gets colder than its
freezing point, but because crystals of the solid haven’t formed, the solution doesn’t
freeze.
Dry your large test tube and accurately weight it to three decimal places. Obtain about 15
mL of cyclohexane and place this in your test tube. Weigh the test tube and the
cyclohexane together to get an accurate determination of the mass of cyclohexane you
will be using today.
Fill a 400 mL beaker with an ice-water mix. Place your thermometer and wire stirrer in
your test tube, and then place the test tube assembly in your ice-water beaker. When the
temperature of the cyclohexane gets lower than 15oC, record the temperature of the
solvent every 15 seconds until the hexane has frozen, about 5-10 minutes.
When the experiment is over, heat the solution back to room temperature, and repeat the
experiment 2 more times.
Experiment 2
Determination of the molar mass of an unknown substance
When a nonvolatile solute is added to a solvent the freezing point is depressed. From
Equation 1 you know that the freezing-point depression is a function of the solvent’s
molal freezing-point depression constant and the solute’s molal concentration.
The molal freezing point depression constant for cyclohexane is 20.0 oC @ kg/mol.
The freezing point depression for the unknown substance will be determined in much the
same way as the freezing point of the pure solvent. A weighed amount of the unknown is
placed in the test tube and completely mixed with the solvent. A cooling curve for the
140
solution is then determined. One change that will be observed is the solution will not
freeze at a single temperature as the solvent did, but will instead freeze over a range of
temperatures. This happens because the pure solvent is freezing on the outside of the tube
while the solute remains in the liquid and gets more and more concentrated as more and
more of the solvent freezes to the tube. Which temperature corresponds to the true
freezing temperature? The initial temperature that is observed before the freezing process
begins. How do you find that temperature? You must fit your data with two straight
lines, one that fits the initial cooling of the solution, and a second that fits the freezing of
the slush mixture. Where these lines intersect is the freezing point of the solution.
Accurately weigh about 0.1 gram of the unknown and add it to your cyclohexane
containing test tube. Determine the cooling curve for this mixture.
Warm the test tube to room temperature, accurately weigh and add another 0.1 gram
sample of the unknown to the test tube and determine the cooling curve of this solution.
Repeat one more time.
Determine the ÄT for the three solutions. From the ÄT of the solution on the molal
freezing-point depression of cyclohexane determine what the molal concentration of the
three solutions must have been. From the molal concentration of the solutions and the
weight of the solute used to make the solution, determine the molecular weight of the
unknown. Remember that you added more and more solute to the solution, so you should
be working with the cumulative weight of solute in each solution. The empirical formula
for the unknown is C3H2Cl.
Don’t forget to make a diagram of your experimental setup in you lab notebook.
141
Name:
Report Sheet
Freezing point Depression
Mass of test tube and cyclohexane______________ (g)
Mass of test tube
______________ (g)
Mass of cyclohexane (the solvent)______________(g)
I. Freezing point of cyclohexane (solvent)
Run I
Freezing point of cyclohexane
________
Average freezing point of cyclohexane:
Run II
________
Run III
________
______________oC
II. Freezing point depression of unknown
Solute Mass data
Mass of unknown added in run 1 ______________g
Mass of solute____________
Mass of unknown added in run 2______________g
Total mass of solute now ____________
Mass of unknown added in run 3 _____________g
Total mass of solute now ____________
142
Freezing point data
Run
Freezing point (oC)
Freezing point depression (ÄT)
1
2
3
Analysis
Kf for cyclohexane _____________________
Molality of solution run 1 ___________________
Molality of solution run 2 ___________________
Molality of solution run 3 ___________________
Show calculation of molality:
Moles of solute present in solvent:
Run 1 _______________
Run 2 ______________
Run 3 _______________
Calculation of moles of solute:
143
Molecular Weight of solute:
Run 1__________
Run 2 __________
Run 3 __________
Calculation of molecular weight:
Average molecular weight_____________________
Molecular Formula of Unknown: ____________
Please attach copies of all Excel graphs used in your analysis
144
Experiment 19.
Rate Law for the Iodine Clock Reaction
Purpose
In this experiment you will study the concentration dependence of the rate of
a reaction.
Background
The rate of a chemical reaction is dependent on a number of physical and chemical
properties of the reactants. In this experiment, we will observe and measure the effect of
concentration of the reacting species on the rate of a reaction. After plotting the data to
determine the reaction order for each reactant, the rate law may be deduced.
The reaction to be studied is the reduction of iodate ion by hydrogen sulfite ion in acid
medium:
3 IO3- (aq) + 8 HSO3- (aq) + 2 H+ (aq) Y 8 HSO4- (aq) + I3- (aq) + H2O(l)
The progress of the reaction will be followed by adding starch to the reaction mixture and
observing the appearance of a blue colored complex which forms between the I3- and the
starch. The blue color will appear when the triiodide ion concentration reaches a certain
level, so the time that it takes to appear is a measure of the rate of the reactions.
To simplify the data analysis, the reaction order for each reactant should be determined
separately. This is done by holding the concentrations of two of the three reactants at
some constant level and varying the concentration of the third. For example, to determine
the reaction order for the iodate ion, you do a series of reactions with different iodate ion
concentrations in each run but with constant hydronium ion and hydrogen sulfite ion
concentrations in all of the runs. Then, in another sequence you vary the hydrogen sulfite
ion concentration at constant iodate ion and hydronium ion concentrations, and so on.
Tabulate your data using as headings "mL potassium iodate", "mL sodium sulfite", “mL
sulfuric acid”, and "time required for reaction.” Convert all of the reactant levels to molar
concentrations in the 100 mL of reacting solution.
145
Procedure
Form teams of two or three. Team members will work together on Part I. Divide labor on
Parts II, III, and IV. If you divide up the runs for a given one of the Parts, be sure to use
the same stock solutions for all of the runs.
Experiment 1
Preparation of solutions
There should be about 1liter of H2O in a round bottom flask at your desk for making the
following solutions.
Weigh out 3.4 g of KIO3 (s) and transfer this to a clean 600 mL beaker. Add 400 mL of
distilled water and swirl until the salt has dissolved. This is the stock iodate ion solution.
Label the beaker.
Weigh out 0.25 g of Na2SO3 (s) and transfer this to a clean Erlenmeyer flask. Add 200 mL
of distilled water and swirl until the salt has dissolved. This is the stock sulfite ion
solution. Label the flask.
Measure out 4.0 mL of 3.0 M H2SO4 (aq) in a 10 mL graduated cylinder and transfer to
another clean Erlenmeyer flask. Add about 200 mL of distilled water. This is the stock
acid solution. Label the flask.
Obtain about 20 mL of the starch solution in a clean 50 mL beaker.
Clean a buret and fill it with your stock iodate ion solution. The buret will be used to
deliver the proper volume of this solution directly into the reaction vessel. Use your large
graduated cylinder to measure the water needed, your medium graduated cylinder to
measure the sulfite ion solution, and use your small graduated cylinder to measure the
acid solution. The starch will be added by filling an eyedropper and squiring the entire
dropper full into the reaction vessel.
Experiment 2
Rate Dependence on Iodate Ion Concentration.
Measure 10.00 mL of the iodate ion solution from your buret into a clean 400 mL beaker
placed on a sheet of white paper. Add (in this order) 75.0 mL distilled water, 5.0 mL of
the acid, and one dropperful of starch solution. Measure 10.00 mL of the sulfite solution
146
into the graduated reserved for that solution. Observe your watch, or the wall clock,
which can be read to seconds, and quickly pour the cylinder of sulfite solution into the
beaker, rapidly mix the contents using a stirring rod, and note the time. Continue stirring
until a blue color appears. Record, to the nearest second, the time required for the blue
color to appear.
Repeat the experiment using 20.0 mL of iodate ion solution, 65.0 mL water, 5.0 mL acid,
one dropper full of starch and 10.0 mL of sulfite ion solution.
Repeat the experiment using 30.0 mL of iodate ion solution, 55.0 mL water, 5.0 mL acid,
one dropper full of starch and 10.0 mL of sulfite ion solution.
Repeat the experiment using 40.0 mL of iodate ion solution, 45.0 mL water, 5.0 mL acid,
one dropper full of starch and 10.0 mL of sulfite ion solution.
Experiment 3
Rate Dependence on Hydrogen Sulfite Ion Concentration.
Repeat the experiment using 20.0 mL of the iodate ion solution, 70.0 mL water, 5.0 mL
acid, one dropper full of starch, and 5.0 mL of sulfite ion solution.
Repeat the experiment using 20.0 mL of the iodate ion solution, 65.0 mL water, 5.0 mL
acid, one dropper full of starch, and 10.0 mL of sulfite ion solution.
Repeat the experiment using 20.0 mL of the iodate ion solution, 60.0 mL water; 5.0 mL
acid, one dropper full of starch, and 15.0 mL of sulfite ion solution.
Repeat the experiment using 20.0 mL of the iodate ion solution, 55.0 mL water, 5.0 mL
acid, one dropper full of starch, and 20.0 mL of sulfite ion solution.
Experiment 4
Rate Dependence on Hydronium Ion Concentration
Repeat the experiment using 20.0 mL iodate ion solution, 67.5 mL water, 2.5mL acid, one
dropper full of starch, and 10.0 mL of sulfite ion solution.
Repeat using 20.0 mL iodate ion solution, 65.0 mL water, 5.0 mL acid, one dropper full
of starch, and 10.0 mL of sulfite ion solution.
147
Repeat using 20.0 mL iodate ion solution, 62.5 mL water, 7.5 mL acid, one dropper full
of starch, and 10.0 mL of sulfite ion solution.
Repeat using 20.0 mL iodate ion solution, 60.0 mL water, 10.0 mL acid, one dropper full
of starch, and 10.0 mL of sulfite ion solution.
Calculations
Each team member should collect raw data from her/his partners and perform the
calculations independently.
The experimental rate law for a reaction is an equation relating the observed reaction rate
to a rate constant k and the known concentrations of the reactants ([IO3-], etc.) raised to
some powers (i.e. reaction orders a, b, c). Since the rate is inversely proportional to the
observed reaction time t, we may write
1/t % rate = k [IO3-]a [HSO3-]b[H+]c
Since we are mainly interested here in determining the reaction orders and not k, we may
rewrite this equation after taking the logarithm as
log (1/t) = constant + a log ([IO3-]) + b log ([HSO3-]) + c log ([H+]).
When only one of the reactant concentrations, such as [IO3-], is varied, the equation
simplifies to the linear form
log (1/t) = constant + a log ([IO3-]),
and we may obtain the value of “a” as the slope in a plot of log (1/t) versus log([IO3-]).
The values of “b” and “c” are found similarly. The advantage of a log-log plot is that the
data should fit to a straight line, which is easier to draw by hand than a curved line fit.
Using your raw data, tabulate your team's log(l/t) and log([X]) values, where [x] refers to
the concentration being varied in each series. Make three separate plots and determine the
reaction orders for each species. Do the experimental reaction orders have integer or noninteger values? Why? Write the overall rate law for the "iodine clock" reaction.
148
Name:
Report Sheet
Rate Law for the Iodine Clock Reaction
I. Preparation of Solutions
starting
material
amount (g or
ml)
volume of
water added
(mL)
reactant
stock
concentrations
(mol/L)
KIO3
Na2SO3
H2SO4
Calculations
Show how your team calculated the concentration of each of the reactants above.
[IO3-]:
[HSO3-]:
[H+]: (Note: Assume that 1 mol of H2SO4 produces 1 mol of H+)
149
II. Rate Dependence on [IO3-]
The amounts of HSO3- and H+ were held constant during this part.
HSO3-: volume (mL): ____________ diluted concentration (M) = ____________
H+ :volume(mL):_____________ diluted concentration (M) = ______________
run #
vol. stock IO3(ml)
log [IO3-]
diluted
[IO3-] (M)
reaction
time (s)
log(1/time)
Calculation
Show how you calculated [IO3-]for run #1:
Plot: Attach a graph of log(1/t) vs log [IO3-]. The slope of this plot is the order of the
reaction with respect to IO3-. (Graph paper may be found on following pages)
IO3- reaction order = _________________
III. Rate dependence on [HSO3-]
The amounts of IO3- and H+ were held constant during this part.
IO3-: volume (mL) = ______________diluted concentration (M) = ____________
H+: volume (mL) = _______________diluted concentration (M)= ___________
run #
vol. HSO3stock (mL)
diluted
[HSO3-] (M)
log [HSO3-]
150
reaction time
(s)
log (1 / t )
Calculation
Show how you calculated [HSO3-] for run #1:
Plot: Attach a graph of log (1/t) vs log [HSO3-]. The slope of this plot is the order of the
reaction with respect to HSO3-.
HSO3- reaction order = ______________
IV. Rate dependence on [H+]
The amounts of IO3- and HSO3- were held constant during this part.
IO3- : volume (mL) = ________________
diluted concentration (M)____________
HSO3- : volume (mL) = ______________
diluted concentration (M)____________
run #
vol. (mL)
stock H+
diluted [H+]
(M)
log [H+]
reaction time
(s)
log (1 / t)
Calculation
Show how you calculated [H+] from run #1:
Plot: Attach a graph of log (1/t) vs log [H+]. The slope of this plot is the order of the
reaction with respect to H+.
H+ reaction order = _________________
V. Rate Law
Summarize your results by writing the overall rate law for the iodine clock
reaction:
rate = _____________________________________
151
Experiment 20.
Reaction Rates
Purpose
In this experiment you will investigate various different factors that influence
the rates of chemical reactions.
Background
Chemical reactions occur as a result of collisions between atoms, molecules, or ions of
the reactants (reacting substances). The rate or a reaction is defined as the change
concentration of a substance per change in time. The units or a reaction rate are often
mol L-1s-1. Many reactions take place in steps, but for most reactions, it is only the
disappearance of the reactants and the appearance of the products that can be observed.
The rate of reaction can be determined experimentally by measuring the concentration of
one of the reacting substances or a product as a function of time. Accurate quantitative
measurements of reaction rates require carefully designed experiments. However,
qualitative observations of a variety of chemical changes are possible. For example; if a
gas is evolved from a solution, one can observe the bubbling to see how fast the reaction
is occurring. If a colored product is formed, one can observe the change in color with a
lapse of time.
We know that combustible substances vary to a great degree in the ease with which they
can be ignited. White phosphorus must be stored under water to prevent contact with the
air; otherwise it would ignite at room temperature. On the other hand, there are many
substances that are very difficult to burn. For a reaction to occur, the molecules of the
reactants must be in physical contact with each other. The rates of reactions are affected
by several conditions that alter the activities of molecules. Let us enumerate the
conditions on which the reactions depend: the chemical and physical nature of the
reactants, the concentration of reactants, the temperature, and the presence of a catalyst
Nature of Reactant
Chemical. Potassium combines with oxygen much more rapidly than does lead under the
same circumstances. Platinum does not combine directly with oxygen at all.
Physical (Surface Area). A large piece of wood will catch fire more slowly than thin
shavings of this same wood. A large piece of coal will burn much easier if it is broken
into small pieces. An increase in the amount of surface exposed increases the reaction
rate, assuming that other conditions are the same.
152
Concentration of Reactants
Since the rate of a reaction depends upon the frequency of molecular collisions, it is
reasonable to believe that substances would burn more rapidly in pure oxygen than they
do in air. In air, only about 21 percent of the molecules are oxygen, hence, when a
substance burns in air, the concentration of oxygen is important. To cite a specific
example, sulfur burns more rapidly in a bottle of oxygen than it does in air.
The same dependence of reaction rates on concentration is true for solutions. For instance
Zn metal will react faster with 6 M HCl (aq) than with 1 M HCl (aq) and with 1 M HCl
(aq) faster than with 0.1 M HCl (aq).
Temperature
As a general rule, the rate of a chemical reaction is approximately doubled by a
temperature increase of 100C. For example, coal is stable in air at room temperature, but
above its ignition temperature it burns rapidly. The rates of oxidation of iron, aluminum,
and lead also take place much more rapidly at elevated temperatures. Furthermore, it is
easy to observe that reaction rates in solutions increase as temperature increases. For
example, consider the reduction reaction of a red solution of iron(III) thiocyanate with
tin(ll) chloride. Assume that two solutions of iron(III) thiocyanate are treated with equal
quantities of tin(ll) chloride at different temperatures, one at 200C and the other at 300C.
The reaction rates may be compared by observing the fading of the red color, and in this
case the fading should be about twice as fast in the warmer solution.
The reaction which we will investigate is that between the thiosulfate ion (S2O32-) and a
strong acid. The net ionic reaction is:
2 H+(aq) + S2O32-(aq) Y S(s) + SO2(g) + H2O
The insoluble sulfur gives the solution a milky appearance. The rate of the appearance of
the white, milky precipitate serves as an indication of the rate at which the reaction
proceeds.
Catalyst
Many reactions are affected by a catalyst or an inhibitor; some are accelerated (catalyst)
and other are retarded (inhibitor). Potassium chlorate decomposes very slowly into
potassium chloride and oxygen at 4000C.
2 KClO3 (s) Y 2 KCl (s) + 3 O2 (g)
When manganese dioxide (catalyst) is added, decomposition occurs at a much lower
temperature.
153
In the synthesis of ammonia, it is necessary to use a catalyst in order for nitrogen to react
with hydrogen.
N2 (g) + 3 H2 (g) X 2 NH3 (g)
The reaction which we will investigate is the decomposition of hydrogen peroxide
according to the reaction:
2 H2O2 Y 2 H2O + O2 (g)
The evolution of the gas, O2, causes bubbles which may serve as an indication of the rate
of reaction.
Procedure
Experiment 1
The Effect of the Nature of the Reactants (Physical: Surface Area)
Place a tin shot in one test tube and an equivalent amount of granular tin in a second test
tube. Add 2 droppers full of 12 M HCl (aq) to each test tube. Keep the concentrated acid
in the hood. Compare the rates of reaction. Explain.
Experiment 2
The Effect of Temperature
Measure 10 mL of 0.050 M sodium thiosulfate solution into each of seven clean test
tubes. Make sure these test tubes have been throughly washed, since traces of acid from
the previous experiments will ruin this experiment. Place one of the test tubes in a 400
mL beaker of cold water (about 150C) such that the water in the beaker level is slightly
above the level of the solution in the test tube. Obtain a watch that can be read to
seconds. Measure 1.0 mL of 0.1 M sulfuric acid in a 10 mL graduated cylinder. Pour this
acid into the thiosulfate solution, tightly cover the test tube with a piece of parafilm, then
quickly turn the test tube upside-down one or two times to mix the acid and the
thiosulfate, and return the test tube to the water bath and start the timer. Watch the
solution closely and note the time at which a faint cloudiness first appears. Record the
elapsed time and temperature of the water in the beaker.
Remove the test tube in which the reaction has occurred, rinse and wipe it dry. Add hot
water to your beaker until the thermometer shows that it is about 5o C warmer than your
previous experiment. Now place another thiosulfate test tube in the beaker and give it a
minute or two to come to the same temperature as the water bath. Now add acid to this
test tube and repeat your experiment at this higher temperature. Record this time and
154
temperature. Repeat this procedure for successive 5 0C rises in temperature, using the
remaining five samples. Tabulate the temperature (T) and elapsed time (t) for each
solution. Estimate each reaction rate as 1/t.
Plot ln(rate) versus 1/t (in Kelvin) and fit a straight line through the points. Determine the
activation energy for the reaction from the slope of the fitted line. How does this result
support the theory that only “activated” molecules are capable of reacting upon collision?
Experiment 3
The Effect of Catalysts
A) Pour 2 droppers full of 3% hydrogen peroxide solution into each of two clean test
tubes. Add a pinch of manganese dioxide to one of the tubes and note what occurs.
Explain.
B) Place 2 droppers full of 3% hydrogen peroxide solution in each of three test tubes. To
the first and second test tubes add 5 drops of 1M iron(III). What happens to the rate of
H2O2 decomposition in the presence of iron(III)? Add 5 drops of 6M sulfuric acid to the
first test tube only. What happens to the appearance of the iron (III)? What happens to the
rate of H2O2 decomposition? Explain.
Experiment 4
The Effect of the Nature of the Reactants (Chemical Reactivity)
A) Place 2 droppers full of 2 M HCl (aq) into one test tube, 2 droppers full of 2 M HOAc
(aq) into a second test tube and 2 droppers full of 2 M HOAcCl (aq) (chloroacetic acid)
into a third test tube. To each of these test tubes add a piece of magnesium ribbon.
Observe the difference in rates of reaction. Explain.
B) Place 2 droppers full of 6M HCl(aq) in each of three test tubes. Carefully sprinkle a
few granules of zinc into one test tube, magnesium into another and tin into the third.
Compare the rates of reaction.
155
Name:
Report Sheet
Reaction Rates
I. Effect of Reactant Size
Explain how the physical shape of a tin sample affects the rate of the reaction with
HCl(aq):
II. Effect of Temperature
Complete the following table with your data for the reaction of S2O32-(aq) with
H+(aq) at seven different temperatures. Plot the natural logarithm of reaction rate (rate =
1 / time) versus the reciprocal of the absolute temperature, fit the data to a straight line,
then determine the slope of this line. Include with this report sheet a print out of the
graph.
Run #
Observed T
(oC)
1/T(K-1)
Observed
reaction time (s)
rate
(1/time)
ln(rate)
1
2
3
4
5
6
7
Find the activation energy, Ea, with appropriate units from your graph. Show how you
determined Ea.
156
III. Effect of a Catalyst
1. Which of the following statements describing the decomposition of hydrogen
peroxide (H2O2) in the presence of manganese dioxide (MnO2) is FALSE:
a) MnO2 remains in the solid state throughout the course of the reaction
b) the solution is colorless, even after gas bubbles have ceased to form
c) MnO2 is consumed during the reaction; more than a small amount is needed
d) the rate of formation of gas bubbles increases when MnO2 is present
2. What happened to the rate of H2O2 decomposition when you added Fe3+? When you
added Fe+3 and acid? Explain.
IV. Effect of the Nature of the Reactants
1. Explain why magnesium ribbons react differently with HCl(aq) CH3COOH(aq), and
ClCH2COOH(aq), even though the acids have the same concentration.
2. Explain why tin, zinc, and magnesium metals have different rates of reaction with the
same acid solution:
157
Experiment 21.
Solubility of Calcium Iodate
Purpose
In this experiment you will determine the solubility constant for calcium
iodate hexahydrate.
Background
This experiment is designed to acquaint the student with the nature and the determination
of solubility products for slightly soluble salts. The salt to be studied is calcium iodate
hexahydrate.
When salt crystals dissolve in a solvent, the energy released as the ions become solvated
more than compensates for the lattice energy which held the ions together in the crystal
before they dissolved. The greater freedom (measured by the quantity T ÄS) enjoyed by
the dissolved ions with their associated solvent molecules helps to keep the ions in the
solution. The balance between these various kinds of energy determines how much of a
salt will dissolve.
An excess of salt crystals in contact with a solution of the salt will eventually result in a
saturated solution. The weight of salt in a given volume of a saturated solution is a
measure of its solubility which for purposes of this experiment, will be expressed in
molarity.
Solubilities of the various salts differ greatly. For instance, a saturated solution of
mercury (II) sulfide contains only about 10-26 mole of each of Hg2+ (aq) and S2- (aq) per
liter. Thus, one would find on average only one mercury (II) ion and one sulfide ion in
every thousand or so liters of solution. On the other hand, calcium chloride is so soluble
that it can dissolve in water it spontaneously removes from moist air.
When a crystal of salt dissolves to form solvated ions, the quantity of positive charge
associated with the positive ions in solution is equal to the quantity of negative charge
associated with negative ions. (The only region of the solution where this may not be
strictly true is at the surface of crystals in contact with a saturated solution. Even so, the
total positive-to-negative charge-ratio in the bulk of the solution is never detectably
different from one.) This is called the principle of electroneutrality and will always apply
to electrolyte solutions.
158
Consider the salt Ca(IO3)2, dissolving in water to form a saturated solution. The reaction
is:
Ca(IO3)2(s)
º Ca2+ (aq) + 2 IO3- (aq)
The symbol “aq” is used to represent water being consumed to hydrate the ions. The
equilibrium constant for the reaction is:
Ksp = [Ca2+] [IO3-]2
The subscript “sp” indicates that K is a special kind of equilibrium constant called a
solubility-product constant. The concentration of Ca(IO3)2 (s) does not appear in the
equilibrium constant expression because it does not vary appreciably as the salt dissolves.
One Ca2+ ion is formed as each unit of Ca(IO3)2(s) that dissolves. Thus:
[Ca2+] = [Ca(IO3)2]dissolved
For every Ca2+ ion there are two IO3- ions appearing as products in solution.
Therefore:
[IO3-] = 2 [Ca2+] = 2 [Ca(IO3)2 (s)]dissolved
Then, Ksp
=[Ca(IO3)2(s)]dissolved
(2 [Ca(IO3)2(s)]dissolved)2
= 4 ([Ca(IO3)2]dissolved)3
To calculate Ksp one need only to analyze a measured volume of saturated solution for
Ca2+ or IO3- . The IO3- concentration can be determined conveniently using the reaction of
this ion with iodide ion according to
IO3- (aq) + 8 I- (aq) + 6 H+ (aq) Y 3 H2O + 3 I3- (aq),
followed by determination of I3- concentration by titration with a standard thiosulfate
solution according to:
I3- (aq) + 2 S2O32- (aq) Y 3 I- (aq) + S4O62- (aq).
159
Procedure
Get about 120 mLs of a saturated solution of calcium iodate in a labeled 250 mL beaker
(The solutions were prepared a few weeks ago. Solid calcium iodate was added to water,
and since then, some of the solid has dissolved until the solution became saturated.) Filter
this through filter paper, which is folded and inserted into your glass funnel, into a clean,
dry 250 mL beaker. Pipet 25.00 mL of the filtrate into a 250 mL Erlenmeyer flask, and
add 2 g of KI crystals, 5 mL of
6 M H2SO4 (aq) and 25 mL of distilled water (note that a reaction takes place in the flask).
Titrate with a standard sodium thiosulfate solution (the standard titrant is available in the
lab; fill a 250 mL beaker about one-half full of this, and do not waste it). Use 2 mL of
starch solution as indicator, added when the solution has become pale yellow which is
just before the endpoint. The endpoint is the disappearance of the blue color of the
starch-I2 solution. Calculate [Ca2+ (aq)] in the saturated Ca(IO3)2 (aq) solution. Repeat the
titration three more times with three more 25.00 mL portions of the filtrate. Calculate Ksp
of Ca(IO3)2 @ 6 H2O. Save the remainder of the filtrate for the qualitative tests.
Qualitative tests on a saturated calcium (II) iodate solution
Many cations form relatively insoluble iodate salts. One of the most insoluble is Pb(IO3)2
whose Ksp is 1.2 x 10-13.
A. Test: To about 1 mL of your Ca(IO3)2 solution add 2 drops of 0.10 M Pb(NO3)2 (aq).
Record your observation in your notebook and on the data sheet. Write the equation for
the reaction.
Calcium ion forms insoluble salts with a few anions other than IO3- , some of which are
less soluble than the iodate salt. These include the carbonate and oxalate ions.
B. Test: To about 1 mL of your Ca(IO3)2 solution add 1 mL of 0.10 M Na2CO3 (aq).
Record your observations in your notebook and on the data sheet. Write the equation for
the reaction.
160
Questions for your Notebook
1.
Write the oxidation half-reaction, reduction half-reaction and the net ionic
equation for the reaction of an aqueous acidic solution of iodate ion with excess iodide
ion.
2.
Write the oxidation half-reaction, reduction half-reaction and the net ionic
equation for the reaction of an aqueous acidic solution of I3- with S2O32-.
161
Names: _________________
_________________
Report Sheet
Solubility of Calcium Iodate
I. Concentration of Thiosulfate Titrant Solution:
Stock [Na2S2O3] (M) = ________________
II. Data and Calculations
Quantity
trial 1
trial 2
Final Buret
Initial Buret
Volume (mL)
Na2S2O3 (aq)
volume (mL)
Ca(IO3)2(aq)
[IO3-] (M)
[Ca2+] (M)
Ksp
Show how you calculated [IO3-] for trial 1:
Show how you calculated [Ca2+] for trial 1:
162
trial 3
trial 4
Show how you calculated Ksp for trial 1
III. Results
Report your average value for Ksp and its standard deviation.
Ksp = _________________________± __________________
(Average)
(std. dev.)
IV. Discussion Questions
1. Write and label the half-reactions for the oxidation of I- (aq) by IO3- (aq) in acidic
solution.
2. Write and label the half-reactions for the reduction of I3- (aq) by S2O32- (aq).
3. Write a balanced net ionic equation describing the reaction between aqueous sodium
carbonate and a saturated calcium iodate solution.
163
4. Ksp = 1.2 x10-13 for the reaction Pb(IO3)2 (s) W Pb2+ (aq) + 2IO3- (aq). Suppose that a
student mixes 1.00 mL of a 1.00 x 10-7 Pb(NO3)2 (aq) solution with 10.0 mL of a 3.00 x
10-3 M Ca(IO3)2 (aq) solution. Will a precipitate form? _____________ Show a
calculation to support your answer.
164
Experiment 22.
Acid-Base Strength of Salts
Purpose
In this experiment you will use a pH meter to examine salt solutions
containing strong and weak acids and bases.
Background
In a previous experiment, we observed the acid-base characteristics of some oxides and
hydroxy compounds. In this experiment we will study another sequence of acids and
bases not only to classify them according to type, but also to determine their relative acid
and base strengths.
One method for experimentally measuring the relative strength of two acids is by reacting
them with the same weak base and determining the extent of the reaction. The greater the
extent of the reaction, the stronger the acid. A similar process can be used to determine
relative base strengths by reacting them with the same weak acid. A convenient weak acid
(or base) for the comparison, in many cases, is water, an amphiprotic material.
Pure water undergoes autohydrolysis as follows:
2 H2O X H3O+ + OHWhen an acid stronger than water is placed in aqueous solution, a hydrolysis reaction can
occur:
HA + H2O X H3O+ + AThe stronger the acid, the more H3O+ that will be yielded in the reaction, and
the more acidic will be the solution. When a base stronger than water is placed in aqueous
solution, the hydrolysis reaction which occurs is:
B + H2O X HB+ + OHThe stronger the base, the more OH- that will be yielded in the reaction, and the more
basic will be the solution.
165
A quantity which is commonly used to evaluate the acidity or basicity of an aqueous
solution is the concentration (in moles/liter) of the hydronium ion. This ion is always
present to some extent in aqueous solutions because of the autohydrolysis reaction. Its
concentration is very high in acidic solutions due to the acid hydrolysis reaction, and very
low in basic solutions due to the base hydrolysis reaction. The hydronium ion
concentration is usually recorded in terms of a quantity called the pH where
pH = -log [H3O+]. We will study this quantity in much greater detail later on in the course
during the consideration of ionic equilibria. It is sufficient at this point to know that a low
pH value is associated with more acidic solutions, and a high pH value is associated with
more basic solutions. Pure water has a pH of 7.00 at room temperature. The following
diagram summarizes the normal pH range found with weak acids and bases in aqueous
solution:
Weakly
Weakly
Very acidic
acidic
Neutral
basic
Very basic
+)))))))))))))0))))))))))))0))))))))))0))))))))))))),
0
3.5
7
10.5
14
In order to use this scale to determine the relative acidic or basic strength of a salt, we
will measure the pH of the solution prepared by dissolving the salt in water. If the
solution is neutral, then probably neither of the ions formed when the salt ionized in
solution has hydrolyzed to any appreciable extent. If the solution is acidic, then one of the
ions is a stronger acid than water, and it has hydrolyzed to form some hydronium ion. The
ions that are stronger acids will hydrolyze more than the weaker ones, producing more
hydronium ions, and making the pH lower. Likewise, if the solution is basic, then one of
the ions is a stronger base than water, and it has hydrolyzed to produce hydroxide ion.
Stronger bases hydrolyze more, yielding higher pH values.
The pH meter is an electrical device designed to measure the hydronium ion
concentration of an analytic solution. In the study of Voltaic cells, it was noted that
electrodes placed in solution would obtain a potential due to the redox properties of the
chemical constituents of that solution. There also exist electrode systems which are
designed so that the potential on the electrode is independent of the solution species with
which it is in contact. These are called reference electrodes.
If two reference electrodes (with potentials E1 and E2) are placed in a solution and their
potential difference (ÄE) is measured with a voltmeter, then:
ÄE = E1 - E2 = ER
thus ER will be constant for a given set of reference electrodes.
166
If one of these electrodes is separated from the analytic solution by means of a membrane
barrier, a new potential, Em, is introduced into the above equation:
ÄE = ER - Em
where Em is the potential arising from the flow of charge carriers through the membrane.
This is caused by the charge flow properties of the membrane material, the difference in
charge buildup between the inside and outside surfaces of the membrane, and the effect of
the analytic solution on the charge of the outside surface of the membrane. In effect, Em is
dependent on the potentials arising at the inside surface (Ei) and the outside surface (Eo)
of the membrane.
Em = Ei - Eo
For a given reference electrode and membrane, Ei is constant, so that:
ÄE = ER - Ei + Eo = ER + Eo
Membranes can be designed so that the charge on the outside surface (and therefore Eo) is
dependent on the concentration of one specific chemical species in the solution into
which it is immersed. Thus, with ER constant, a change in the meter reading (ÄE)
indicates a change in Eo which is proportional to the concentration of the desired species.
An electrode which is sensitive to hydronium ion contains a glass membrane of specific
chemical composition which interacts selectively with the [H3O+] in the analytic solution.
The potential on the outside surface is related to the pH in the following manner:
Eo = 0.0591 pH
so that:
ÄE = ER + 0.0591 pH
This is an equation of the form y = mx + b, in which the meter reading is directly
proportional to the pH, with an intercept of ER. The pH meter is a voltmeter whose
voltage scale has been replaced with a scale which reads directly in pH units. The
intercept (ER) is initially set by calibrating the instrument with a buffer of known pH. The
pH of any solution can then be read directly from the scale.
167
Procedure
Use of the pH meter
Certain precautions must be observed when operating the pH meter. First, since the tip of
the electrode has a fragile glass bulb, be careful not to bump it against the beaker or hit it
with a stirring rod. Second, do not expose the electrode to air for more than 10 seconds;
transfer it reasonably quickly from one solution to another, and store in water when not in
use. Third, the electrode only measures pH accurately between values of 2 and 12, so it
should not be used in strongly acidic or basic solutions. Fourth, the electrode should be
rinsed off with distilled water during the transfer between solutions, so that samples are
not contaminated by drops of a previous sample.
The use of the computer based pH system will be demonstrated by your instructor. Please
take note of all precautions. The interfaces cost $200 each, and the electrodes cost $50.
BE CAREFUL!
The electrode should be stored in the plastic holder with a pH 4 buffer solution. To make
a pH measurement, remove the electrode, rinse with distilled water from your wash
bottle, blot the electrode gently with a kimwipe, and carefully put it in the solution to be
measured. Make sure it is properly submerged (Check with instructor if unsure). Find the
window on the computer display that gives the pH reading of your electrode. After
immersing your electrode in your solution, wait a minute of two for the pH reading to
stabilize. Gently swirl or stir the solution once or twice during this minute. After the
measurement remove the electrode, and rinse before immersion in the next solution or the
holder.
In order to ensure accurate measurement of the salts, it is imperative that all your
glassware used be immaculately clean. Tiny traces of acid or base leftover from previous
experiments can severely distort a measurement.
168
A. Preparation of the Solvent
In order to obtain proper measurements with the electrode, it is necessary to use water in
which a neutral salt has been dissolved, rather than pure distilled water. Weigh out
approximately 2.5 g of KCl (s), and dissolve in approximately 500 mL distilled water in a
600 mL beaker. Use this solution whenever reference is made to your “solvent.”
B. Measurement of Salt Solutions
1. Place 30 mL of the solvent in a 100 mL beaker and measure the pH. Record this initial
pH. This ought to be a neutral solution, but the pH may be slightly due to CO2 (g) which
dissolves from the air to form carbonic acid, H2CO3 (aq). Without removing the pH
electrode add a large spatula tips of (a) NH4Cl (s) and stir until it dissolves. Record the
final pH. Did the pH go up or down? Repeat the procedure with a fresh sample of
solvent and NaC2H3O2 (sodium acetate).
2. Repeat, using (a) NaNO3 (s), (b) Mg(NO3)2 (s), and (c) Al(NO3)3 (s).
3. Repeat, using (a) NaClO4 (s), (b) Na2SO4 (s) and (c) Na3PO4 (s).
4. Repeat using (a) NaH2PO4 (s), and (b) Na2HPO4 (s). Compare these two to Na3PO4 (s)
in part 3 (above).
Questions for your Notebook
1. For parts 2-5, construct the following table.
Salt
Initial pH
of solution
Final pH
Acid or
Ion
of solution Base
Responsible
Conjugate form Overall
of ion
Strength
2. For each of parts B.2, B.3 & B.4, rank the substances by extent of hydrolysis
(pH); e.g. A > B > C. For each part give an explanation of the differences in extent
of hydrolysis based on charge density.
169
C. Comparison with acidic and basic oxides
1. Put a fresh 30 mL portion of solvent into each of two beakers. Into one beaker dissolve
a spatula tip of (a) CaO (s) and into the other beaker dissolve a spatula tip of (b) CrO3 (s).
Measure the pH of each.
2. Into a fresh 30 ml portion of the solvent, use a straw to gently blow 3 breaths of air.
Measure the pH. (The air from your lungs should be a good source of CO2. If you hold
your breath first you will have even more CO2 in your lungs, but be careful - don’t pass
out!)
Questions for your Notebook
1. How do the pH's of the solutions of CaO and CrO3 compare to those for the salts of
Section B?
2. Did the CO2 from your lungs have any effect on the pH of the solution?
3. In general, should the pH’s of ionic oxides be >7 or <7? In general, should the pH’s of
covalent oxides be >7 or <7. Did you observe any exception to these general rules? If
you did observe an exception, can you explain why the exception occurs?
D. Acid-Base Reactions and Conjugate Pairs
1. Place 30 mL of 0.10 M NH3 (aq) in a 100 mL beaker and 30 mL of
0.10 M NH4Cl (aq) in a second beaker. Measure the pH of each solution and
determine which is acidic and which is basic.
To the acidic solution, add 6 M NaOH (aq) dropwise, measuring and
recording the pH after the addition of each drop. Continue adding NaOH (aq)
until the pH rises to a reading of about 12.
To the basic solution, add 6 M HCl (aq) dropwise, measuring and recording the pH after
the addition of each drop. Continue adding the HCl (aq) until the pH drops to a reading of
about 2.
2. Place 30 mL portions of your solvent (KCl in H2O) into each of two 100 mL beakers
and repeat the above dropwise additions of HCl (aq) into one, and NaOH (aq) into the
other.
170
Questions for your Notebook
1. For the two cases in section D.1, make a rough graph of pH vs. drop number. Explain
the shapes of the curves in terms of reactions taking place.
2. Compare the results of Section D.2 to those of Section D.1 and explain the difference.
171
Name:
Report Sheet
Acid-Base Strength of Salts
I. pH of Salt Solutions
salt
Initial
pH
Final
pH
acid, base
or neutral
Ion
responsible
conjugate form of
acidic/basic ion
KCl -‘solvent’
NH4Cl
NaCH3CO2
NaNO3
Mg(NO3)2
Al(NO3)3
NaClO4
Na2SO4
Na3PO4
Na2HPO4
NaH2PO4
1. Which of the following cations has the least positive charge density? {Na+, Mg2+, A13+}
Which one is least acidic? {Na+, Mg2+, A13+}
2. Which of the following anions is most basic?
{CH3CO2-, NO3-, ClO4-, SO42-, PO43-, HPO42-, H2PO4-}
172
II. pH of Oxides
1. Write the net ionic equation describing the reaction between CaO (s) and H2O (l):
2. Write the net ionic equation describing the reaction between CO2(g) and H2O(l):
3. The pH CaO solution is {greater than, less than, equal to} the pH of the CO2 solution.
III. Acid-Base Reactions
1. Write the net ionic equation for the reaction between NH4Cl (aq) and NaOH (aq):
2. Write the net ionic equation for the reaction between NH3 (aq) and HCl (aq):
3. Explain why the pH at the stoichiometric point is not the same for the two reactions
above.
4. Predict the pH (± 1 unit) at the stoichiometric point in the reaction
between
a)
CH3COOH (aq) and NaOH (aq)
pH.
b)
Na3PO4 (aq) and HCl (aq)
pH.
173
Experiment 23.
Buffers and Potentiometric Titrations
Purpose
In this experiment you will use a pH meter to study buffer action of weak
acid/base solutions.
Background
A buffer solution is composed of a weak acid and its salt or a weak base and its salt. Such
solutions can absorb moderate amounts of strong acid or base with only small changes in
pH. For example, in an acid buffer, the salt is able to react with added acid to minimize
pH changes. The weak acid has the same effect on added base. As long as the
concentration of the weak acid or its salt is not exceeded by the added base or acid the
solution will remain buffered. However, a buffer is most effective when the
concentrations of weak acid and its salt are equal. These statements also apply to
solutions of a weak base and its salt.
Buffers play a significant role in chemistry where they may be used to maintain a pH
during a reaction. Buffers also play an important role in the chemistry of living systems.
Such systems are very sensitive to pH changes. One of the most important buffers in
living systems is the CO2:HCO3-:CO32- buffer system. For the above reason, buffers are
also important in many drugs.
It is interesting to compare a buffered system with a non-buffered system of the same pH.
This will be done in this experiment using a H2PO4-:HPO42- buffer of around pH 7 and
comparing it to water.
Procedure
I. Preparation and Properties of a Buffer
1. Obtain about 30 mL of 0.10 M NaH2PO4 solution in a 250 mL beaker and about
30 mL of 0.10 M Na2HPO4 solution in a second 250 mL beaker. Measure the pH
of each solution. Mix these two solutions to prepare a buffer solution and measure
its pH. Also measure the pH of distilled water.
174
2. Place 20 mL of the buffer solution from 1 in a 100 mL beaker and 20 mL of
distilled water in a second 100 mL beaker. Add 1 mL of 0.1 M HCl (aq) to each
beaker and measure the pH’s.
3. Place 20 mL of the buffer solution from 1 in a 100 mL beaker and 20 mL of
distilled water in a second 100 mL beaker. Add 1 mL of 0.1 M NaOH (aq) to each
beaker and measure the pH’s.
4. To study the effect of adding an excessive amount of strong acid to a buffer, add
5 mL of 0.1 M HCl (aq) to the remaining 20 mls of buffer solution from 1. Measure the
pH. Add an additional 5 mL of 0.1 M HCl (aq) to the buffer solution and measure the
pH. Finally add an additional 10 mL of 0.1 M HCl (aq) to the buffer solution and
measure the pH. At what point does the buffer finally stop buffering?
5. Interpret the above results in terms of your knowledge of buffers.
II. Potentiometric Titration
A. Using at 25.00 mL volumetric pipet, deliver 50.00 mL (two 25.00 mL aliquots) of
approximately 0.1 M hydrochloric acid into a 250 mL beaker. Fill a buret with
standardized NaOH (aq) (approx. 0.2 M). Place the pH electrode in the acid solution and
position the buret so that direct additions can be made to the beaker from the
burette. Record the initial pH of the acid solution. Add 1.00 mL of base, swirl the
solution to ensure complete reaction and record the pH reading. Continue adding
1.00 mL portions, swirling, and taking pH readings until the pH begins to rise
more quickly (this should require a total of 23 to 24 mLs of titrant). Add titrant in
0.20 mL portions as the pH is changing rapidly, and continue this until the pH
starts to level off at some high value. Finish the titration with several 1 to 2 mL
additions. In your notebook make a data table with mLs of base in one column and pH in
a second column
Plot this data on graph using excel and attach this graph to your report sheet. Indicate on
this graph where the equivalence (stoichiometric) point is. Calculate the actual molarity
of the hydrochloric acid from the given concentration of the NaOH (aq) and the number
of mLs it takes to get to the equivalence point.
B. Repeat the procedure in A above using 50.00 mL of approximately 0.1 M acetic
acid in place of the hydrochloric acid. Make a data table and plot the data in excel as
above. On the graph indicate both the equivalence point, and the point where pH=pKa .
Also calculate the concentration of acetic acid.
175
Name:
Report Sheet
Buffers and Potentiometric Titration
I. Buffers
Solution
Measured pH
0.10 M NaH2PO4
0.10 M Na2HPO4
buffer
distilled water
distilled water + 1 mL acid
buffer + 1 mL acid
distilled water + 1 mL base
buffer + 1 mL base
Write the net ionic equation for the reaction between HCl (aq) and the buffer solution.
Write the net ionic equation for the reaction between NaOH (aq) and the buffer solution.
II. Potentiometric Titration
1. Plot the titration curves for the two reactions. On each graph, indicate the pH and the
volume of base added at the equivalence point. If done on a computer, include a print out
of your graph with this report.
2. Calculate the concentration of hydrochloric acid used in the first titration.
[HCl] =
mol/L
3. Show calculation:
176
4. Calculate the concentration of acetic acid used in the second titration.
[CH3COOH] =
mol/L
5. Show calculation:
6. Calculate Ka for acetic acid
Ka =
7. Show Calculation:
8. Consider a buffer solution formed by mixing 0.100 mol CH3COOH and 0.100 mol
CH3COOK into 1.00 L water.
a).
Use your data to estimate the pH of this buffer solution. pH =
b).
Write the net ionic equation that describes the reaction that would occur if
HCl (aq) was added to this acetate buffer:
Attach plots of your two titration curves.
177
Experiment 24.
Structures of Organic Molecules
Purpose: In this experiment you will:
1.
build models of organic and biochemical molecules
2.
examine different isomers of compounds with the same empirical formula
3.
examine examples of different organic compounds
Background
The structures of molecules and solids are largely responsible for their physical and
chemical behaviors. Compounds with similar empirical formulas may have very different
properties because the positions of the atoms and the electronic charge distributions may
be very different. Therefore, chemists often write pseudo-structural formulas for a
molecule to indicate precisely how the atoms are connected. For example, we often write
CH3COOH instead of C2H4O2 for acetic acid. A physical model provides even more
details about the spatial relationship of the atoms in a compound. A model may allow us
to view the overall 3-D shape (spherical, flat, rod-like, etc.) of a molecule, to appreciate
how flexible it is, or to gauge the distance between atoms that are not directly bonded to
each other. Models also enable us to predict how entire molecules interact with each other
in a condensed phase. In any event, we must not take the meaning of models too literally,
because models never provide a perfect description of real compounds.
There are several different types of pieces in your plastic model kits. They are color
coded and, in some cases, labeled according to the hybridization (sp, sp2, sp3) of the
center atom. The black pieces usually represent carbon, while red or blue pieces represent
other atoms such as oxygen, sulfur, or nitrogen. Special pieces are available to show
double and triple bonds. Hydrogen atoms are usually indicated by simply leaving the end
of the model piece unconnected. Other terminal atoms in a structure, such as halogens,
may be shown using colored balls. Lone pairs of electrons are often not shown in a
model, but they must be taken into account when establishing the geometry around a
central atom. Complex molecules containing multiple center atoms are built by
recognizing the shape of each center atom, assembling that center, and then connecting
the centers together according to a Lewis structure or drawing of the molecule.
When a molecular model is built, the free rotation about single bonds may allow the
model to be twisted into a variety of conformations. If single bonds link atoms into a ring,
the conformations are limited because the bond angles at each center are fixed. Multiple
bonds are rigid and fix the shape of that part of a molecule. Therefore, atoms bound to a
center with a double bond may be connected in different,
non-interconvertable ways. Molecules that differ only by this kind of connectivity are
178
called geometrical isomers. The shape of a molecule and the types of atoms it contains
determines whether that molecule is polar or non-polar. In a large molecule, it is often
possible to recognize a smaller unit containing a group of atoms bonded in a specific way.
The sub-unit causes a particular chemical behavior of that molecule, so the set of atoms is
called a functional group. Very large “macromolecules” often contain repeating patterns of
simple units; in polymers, these units are called monomers.
Procedure
Work in teams of two or three. Construct models of each of the molecules and solids
listed on the report sheet. Have your instructor check your models and answer the
corresponding questions at each step before you proceed to the next set.
Use caution when smelling the compounds that are stored in the fume hood as some of
them are extremely powerful.
MODEL STRUCTURES YOU WILL BUILD
1. Alkanes: These are simple hydrocarbons with sp3-hybridized carbon atoms. Alkanes
may be gases, liquids, or solids. They are often used as fuels.
2. Alkenes: Carbon-carbon double bonds (sp2-hybridized carbon) do not rotate.
Geometrical isomers are possible. Alkenes, such as ethylene (H2C=CH2 ) and vinyl
chloride (ClHC=CH2 ) may be polymerized.
3. Rings: Hydrocarbons may form ring compounds with either sp3 or sp2 carbons. The
differing shapes and flexibility of these compounds gives them very different physical
and chemical properties.
4. Functional groups: Alcohols (-OH), carboxylic acids (-COOH), esters (-COO-),
amines (-NH2), and aldehydes (-CHO) are distinguished by the presence of a
characteristic group of bonded atoms (-XX).
5. The wide world of molecules: Some very interesting molecules are found in nature
(especially in plants). Other molecules are synthesized for medicine or as explosives (e.g.
trinitrotoluene, “TNT”).
6. Biochemicals: Biological molecules have highly specific functions in living
organisms. Small changes in the shape of a molecule may greatly change its activity.
179
SAMPLES TO SMELL
ethanol
ethyl acetate
acetic acid
vanillin
benzaldehydediethylamine
naphthalene
amylacetate
MODELS TO VIEW
testosterone
guanine
retinal
cytosine
deoxyribose
thymine
180
adenine
Names:_______________
______________
Report Sheet
Structures of Organic Molecules
Have your instructor initial next to each of the structures listed below to confirm that your
model correct. After completing each numbered set of structures, answer the question
with the corresponding number on the next page.
I. MODEL STRUCTURES TO BUILD
1. Alkanes (Interchapter F)
methane (CH4 ) ______ ethane (C2H6 )___________
butane (C4H10) _______ (2 structural isomers)
pentane (C5H12) __________ (3 structural isomers)
propane (C3H8 )_________
2. Alkenes and alkynes (Interchapter G)
ethene (C2H4)_________
2-pentyne _______________
cis-2-butene _________ trans-2-butene _______________(G.2)
cis-1,2-dichloroethene _____________ trans-1,2-dichloroethene _____________(G.1)
3. Cyclic hydrocarbons (Interchapter H)
cyclohexane C6H12 (p.1004) ________ Boat and chair forms
benzene C6H6 (H.1)____________
naphthalene* C10H8, (H.5) ________
4. Different Functional Groups (Interchapter P)
ethanol* C2H5OH (P.1)_________
acetic acid* CH3COOH____________
benzaldehyde* (p. 1014) ________
ethyl acetate* (R.4) _________
diethylamine* ((C2H5)2NH)________
amyl acetate*(p1015) ____________
5. Other interesting molecules
Build these two
TNT (H.4)______________
1,4-dibromobenzene (H.4) _________
Write these structures in your notebook and identify the functional groups
vanillin* (p. 1014) acetylsalicylic acid (p.1015)___________
181
6. Biological Molecules
Build: alanine (T.1) ___________
View premade models :
retinal _________ testosterone _________
deoxyribose (T.7) ________
Adenine _______ Cytosine________ Thymine________ Guanine_______ (T.7)
*
Sample available to view, or smell.
QUESTIONS (numbers correspond to the model structures you built)
1. These hydrocarbon compounds are called alkanes. What is the smallest alkane having
more than one possible structure (isomer)? _____________
How many structural isomers can you find for pentane? __________________
2. Which one of these six compounds has a non-zero dipole moment (i.e. is polar)
_____________________
3. Which one of these three compounds is not flat? __________________
Naphthalene is the smallest of a class of compounds called polycyclic aromatic
hydrocarbons (PAHs), which concern some environmental chemists. Naphthalene has a
very characteristic odor - of what? __________
4. Each of these compounds contains a basic functional group attached to a hydrocarbon
fragment, which may be denoted as R.
Which one is an alcohol (general formula R-OH)?
What is the general
formula for an aldehyde? _____________
Which compound smells
nasty/fishy/ammonia? _________________
Which compound smells fruity?
____________________
5. Acetylsalicylic acid is a popular, well established drug; what is its common name?
____________________________
6. Match each these compounds with one of the following applications:
a sugar_______________
photochemical receptor_____________
an amino acid _____________ male sex hormone ___________
182
Appendix 1.
Data Analysis
Chemistry is a quantitative science: wherever possible it defines its concepts more
sharply by associating numbers with them. Properties associated with numbers are called
quantities and the process of obtaining this number is called measurement. Measurement
of a quantity consists of comparing that quantity with a quantity of the same type. The
quantity used for comparison is a unit.
All quantities take the form
Quantity = Number x Unit.
Thus, we might report that a length = 5 cm. The number (the magnitude of the quantity) is
5 and the unit is the centimeter. Pure numbers are quantifies whose unit is 1; they are also
referred to as dimensionless quantities. Except for pure numbers, quantities without units
are meaningless. A length reported as 5 gives us no information about the unit; it could be
cm, m or some other unit Include appropriate units at all stages of your calculations.
Errors in Measurements
Attempts to measure a quantity, to get its “true value," generally yield values associated
with errors. These errors are of two types, systematic errors and random errors i.e.,
True value = Measured value + Systematic error + Random error
Both these errors have signs, i.e., they can be either positive or negative.
!
Systematic errors can be distinguished from random errors in that they always
have the same value both in magnitude and sign. They arise from the use of poor
quality, incorrectly graduated instruments, bad experimental technique, use of the
wrong experimental conditions etc. They can be avoided by correcting the causes
that lead to them, e.g., calibrating the instruments, improving technique, etc.
!
Random errors, unlike systematic errors, vary from measurement to measurement,
both in sign and magnitude. This unpredictability of random error is a
characteristic feature. According to the theory of probability we can correct for
random error by taking many measurements and taking the average: it is very
likely that there are as many positive errors as negative errors of the same
magnitude so that they cancel out on taking the average. Random errors arise due
183
to uncontrolled or “uncontrollable” factors like fluctuations in temperature and
other experimental conditions, random human errors etc. We can attempt to
minimize them but they cannot be altogether avoided.
To make these concepts more definite we need to define some new quantities. The
average
of a series of n measurements Y1, Y2, Y3, ... Yn of the same quantity is defined
by
where the G indicates summation from i = 1 to i = n. The average
is our best estimate
of the true value free of random error. However, because we calculate
from the
erroneous individual measurements Yi, we cannot be certain that
has completely
eliminated random error. The degree to which we can be certain of
is given by the
standard deviation ó defined by
where again the summation is from i = 1 to i = n. The more the individual measurements
deviate from
(regardless of sign), the greater is the value of ó: the standard deviation
is a measure of the spread of the measurements. The smaller the value of ó the more
certain we can be that the calculated average is free of effects of random error. The
technical term for spread is precision: more precise measurements have smaller values of
ó and the individual measurements deviate less from the average value. We note from the
definition that we can decrease the value of ó by taking more measurements. We have to
strike a balance between reducing ó and not spending too long on the measurement. In
general, we take about three or four measurements. In addition to taking replicate
measurements, it is also important to reduce random variation or “fluctuations” in
experimental conditions.
The average
by itself is not completely informative. We also have to give the standard
deviation a of the series of measurements to indicate its precision. We usually quote the
experimental estimate of value of Y, Yexptl as
184
Note that taking averages reduces the effect of random errors but does not correct for
systematic errors. The error of the average value is defined as
and
where the vertical bars , , signify the absolute values, i.e., without sign. Thus, , 2 , = 2 and
, -2 , = 2.
Measurements with a small value of error are accurate measurements while
measurements with small values of the standard deviation ó are precise measurements.
Most often we cannot calculate errors because we do not know the “true value.” An
important objective in experimentation is to get accurate values by (1) reducing
systematic errors and (2) reducing uncertainty in our average by increasing the precision.
We illustrate the meaning of accuracy and precision by considering the following
problem: A mass whose "true" value is known to be 2.654 g was determined by three
workers who obtained the following values:
Set A
2.652 g
2.658 g
2.654g
average = 2.655 g
%error = 0.04%
Std. deviation = 0.003 g
Sample Calculation
Set B
2.620 g
2.621 g
2.619g
2.620 g
1.3%
0.001 g
Set C
2.660 g
2.680 g
2.640g
2.660 g
0.2%
0.02 g
We have Mass true = 2.654 g and n = 3 for all three sets.
185
1. Average for Set A: Mass =
= 2.655 g
2. % Error for Set A: % error =
= 0.04%
3.Standard deviation for Set A: Using the formula
= 0.003g
Accuracy is a measure of the closeness of the average value to the true value and is
measured by the error or % error. Clearly set A is the most accurate while set B is the
least accurate. Precision is a measure of the closeness of the individual measurements to
the average value and is measured by the standard deviation. Set B is most precise while
C has the least precise measurements. How is it that B has precise measurements but is
still the least accurate? Clearly B has some systematic errors present in his measurement.
When the average was calculated for set A the calculator display read 2.654666667, and
the standard deviation was (3.082...) × 10-3. This means that the results of A are uncertain
by at least 0.003 g or 3 mg; information about errors in the 4th, 5th..etc. decimal places is
not particularly useful. We therefore
!
quote standard deviations and errors to only 1 figure, or at most 2 figures
Thus for A, the standard deviation may be reported as 3 mg or 0.003 g (1 figure), or if
you wanted to carry two figures, 3.1 mg or 0.0031 g (by “rounding-off”). Since the
uncertainty is already in the mg range or in the 3rd decimal place if we give the mass in
grams, it is meaningless to quote, in our average; figures beyond the mg range. We
therefore “roundoff” the average to a mg, i.e., to the third decimal place when giving
mass in grams: 2.654 g. If we wished to give the standard deviation to 2 figures, we
would write the experimental value of A as 2.6547 ± 0.0031 g. In set C the standard
deviation was 0.02 g, i.e., the error is in the 2nd decimal place. We must therefore not
quote, in the average, figures beyond the 2nd decimal place: the correct reporting of C’s
186
value is 2.66 ± 0.02 g. If we had been reporting the masses in mg, the correct way of
reporting the values would be:
A: 2655 ± 3mg
B: 2620 ± 1 mg
C: 2660 ± 20 mg [more comments on this presently]
Significant Figures
When we report a value we are attempting to give information about the size of the
quantity measured. The figures in our-reported value are significant to the extent that they
do this; any figures which convey little or no information about the size of the quantity
would be meaningless or insignificant. We use the following rules to include only those
figures (significant figures) which convey useful information about the size of the
quantity:
!
First calculate the standard deviation, usually to one figure or at most two figures,
never more than two figures. We will use one figure.
!
Figures which correspond to sizes smaller than the standard deviation are
insignificant and should not be included.
!
Round off the average to correspond to the size of the standard deviation.
For set A we found the standard deviation to be 0.003082...g; on rounding-off to one
figure this became 0.003 g. This figure of 3:is in the 3rd decimal place and consequently
we should roundoff the average of 2.654666667 g to the third decimal place as 2.665 g.
There are four significant figures in this result, 2,6,6, and 5. The uncertainty is highest in
the last figure, viz., 5 so it is called the least significant figure; the most significant figure
is 2.
When we do not quote standard deviations, the general convention is that the last quoted
figure is in error by about 1 to 3. Thus in 24.56 mL the last figure is 6 in the 2nd decimal
place, implying an error of about 1 to 3 in the 2nd decimal place, i.e., about 0.01 to 0.03
mL.
In a value like 20.30 mL the error is about 1 in the 2nd decimal place and since all four
figures signify sizes equal to or greater than this error, they are all significant: there are
four significant figures in 20.30 mL. To indicate that the error was in the first decimal
place we would have to report the value as 20.3 mL and the number of significant figures
would be three. Changing the unit of a quantity is a mathematical, rather than an
experimental, operation; it cannot change the number of significant figures. Thus 20.30
187
mL can also be written as 0.02030 L but this still has only four significant figures: the two
zeros in front, since they do not have non-zero figures to their left, do not signify any
meaningful size. We thus have the rule:
!
Zeros with no non-zero figures to their left are not significant
!
Zeros with non-zero figures to their left are significant
Thus we have
Measurement # of Sig. Figs
1.008 cm
4(1,0,0,8)
0.106 m
3(1,0,6)
303.01 cm
5(3,0,3,0,1)
0.00067 g
2 (6, 7)
Avoiding ambiguity in conversion of units
The quantity 14.8 m has three significant figures. If we express this in cm we would tend
to write 1480 cm, wrongly implying four significant figures. To avoid this error use
scientific notation and write 1.48 x 103 cm; this now implies, correctly, three significant
figures. Similarly we earlier saw that for C the average was 2.66 g and the standard
deviation was 0.02 g. The correct way to convert this to mg would be to write (2.66 ±
0.02) x 103 mg rather than 2660 ± 20 mg.
Getting the best out of the measuring instruments
The uncertainty in measurements cannot be smaller than the smallest measurements
(sensitivity) the instruments used can make. However, with good technique it is possible
to get precisions approaching the instruments' sensitivities. The analytical balance has a
sensitivity of 0.0001 g and with care we can make weighings whose precision approaches
this. Consequently it is good practice to read the instruments to values consistent with
their sensitivities. When using the analytical balance we would record the mass to the 4th
decimal place in grams. Similarly when using a 10 mL pipet we would record the volume
as 10.00 mL, since with care, we can limit errors to a fraction of a drop ( . 0.05 mL); to
quote the volume as just 10 mL would imply only 2 significant figures, as opposed to the
4 significant figures in 10.00 mL. Burets have graduations 0.1 mL apart, but it is possible
to estimate by eye between the graduations; consequently buret readings can be quoted to
the 2nd decimal place in mL, e.g. 26.43 mL or 24.50 mL. In summary:
!
record weighings on the analytical balance to the 4th decimal place in grams.
188
!
record buret readings to the 2nd decimal place in mL
!
record pipet volumes to the 2nd decimal place in mL
Propagation of errors
When quantities that are in error are combined by mathematical operations (addition,
subtraction, multiplication, division, taking logarithms, exponentiation, etc.) the error in
the result will obviously depend on the errors of the quantities used in the calculation and
the type of calculation. The following are approximate rules for determining the error and
the number of significant figures to be retained for a few simple operations:
!
Do operations in brackets first, then multiplication/division and finally
addition/subtraction. Within brackets, do the operations in this order.
!
In addition and subtraction, the error cannot be smaller than the largest error in the
quantities added or subtracted. First determine the smallest number of significant
figures after the decimal point in the quantities added or subtracted. The result will
have the same number of significant figures after the decimal point. In addition
and subtraction we first decide the number of significant figures after the decimal
point and allow this to determine the total number of significant figures.
!
In multiplication and division determine the smallest number of significant figures
in the quantities used. The final result will have the same number of significant
figures. In multiplication. and division we first decide the (total) number of
significant figures and then determine the error.
189
Examples
Addition
19.82 g
1.981 g
+ 0.1425 g
21.9435 g
First determine the quantity with the smallest number of significant figures after the
decimal point. This is 19.82 g with 2 significant figures after the decimal point. Therefore
the result will also have two significant figures after the decimal point. The result
(21.9435 g) is rounded-off to 21.94 g.
Multiplication
0.621 m x 0.21 m = 0.13041 m2
The factor with the smallest number of significant figures is 0.21 m, which has 2
significant figures. Therefore the result should be quoted to 2 significant figures. The
calculator result of 0.13041 m2 is rounded off to 0.13 m2.
Division
12.841 g ÷ 1.25 cm3 = 10.2728 g/cm3
The number with the smallest number of significant figures is 1.25 cm3, and has 3
significant figures. The final result should be quoted to three significant figures. The
calculator result of 10.2728 g/cm3 is rounded off to 10.3 g/cm3. This has 3 significant
figures but the error is in the first decimal place.
Some factors are exact and have no error because they arise from theory, not experiment
In Diameter = 2 x radius, 2 comes from theory and is exact: it really stands for
2.0000000... and should not be interpreted as having only one significant figure. In a
conversion like kg = 1000 g, 1000 has no associated errors. The number 1.48 x 10 has 3
significant figures if we treat 10 as exact.
Rounding off
In rounding off, the last significant figure is increased by 1 if the figure that follows it is 5
or greater. Thus 0.16419 rounded to 2 sig figs is 0.16 since the following numeral 4 is less
than 5 while 0.17682 rounded off to 2 sig figs is 0.18, since 6 >5.
190
Using Excel to Graph Data
(Office 2013 version)
Objectives:
<
<
learn to use Excel to graph data
learn how to use the trendline function to plot a line of best fit from a set of
data points
Graphing is an important method of analyzing data gathered from experiments. Graphing can
make it easier to see trends in sets of numbers visually for comparison; or in the case of an
experiment or assay that involves plotting a line of best fit for your data points, it is
indispensable. Today we will use the Chart Wizard in Excel to plot sets of data points, then plot
a line of best fit for the data, and finally have the equation for the line printed out on the plot (the
line equation is written as y = mx + b, where m is the slope of the line and b is the y-intercept).
Exercise I of this lab will lead you through the process of setting up a graph step by step. In
Exercise II you will set up a graph for a different data set and learn how to designate data series
and apply trendlines to more than one series.
Exercise I.
Step 1
find a computer with Microsoft Excel installed on it
Step 2
select Excel and open the program
Step 3
click on the ‘Blank Workbook’, the top left template
Step 4
under Page Layout select ‘Size’ and click on ‘Letter’ 8.5" x 11"
under Page Layout select ‘Orientation’ and click on ‘Portrait’
under Page Layout select ‘Margins’ and click on ‘Normal’
under Page Layout click on the lower right hand corner of the Page Layout box to
bring up more options
Under the ‘Header/Footer’ tab
Click on the ‘Custom Header’ button
Click in the Center Section box and type
“Exercise I” and your name
Click on ‘OK’ button
Click on the ‘OK’ button
Note: you could have set the margins and the orientation in the ‘Page Setup’ box
as well!
191
Step 5
Enter the following column titles and data into columns A & B on the Excel
spreadsheet. You can pull the columns wider by placing the mouse pointer on the
column line and clicking and holding the left button while you drag the column
wider. (Note: Excel will by default use data in column A for X values and data
from column B for Y values in a scatter plot. And if you have data in columns
A, B, C, & D highlighted it will plot three series using A as the x value for all
three and B, C, and D as y values)
Data for Enthalpy of
Vaporization of Water
1/T (K -1)
ln (Pvapor)
.00283
-.799
.00286
-.941
.00288
-1.050
.00291
-1.204
.00293
-1.309
.00296
-1.470
.00298
-1.609
.00301
-1.772
.00304
-1.897
.00307
-2.040
.00310
-2.207
Here is how you get the superscripts and subscripts:
First type in the characters ‘1/T (K-1)’
See how the characters are typed in both in the cell and in the long skinny
box above?
Use you mouse to highlight the characters ‘-1' in the upper box
Now right click in the highlighted area
Left click on ‘Format Cells’ in the drop down menu
Left click on the Superscript box
Click on ‘OK’
You should now have 1/T (K-1) displayed on your spreadsheet
Follow a similar procedure to make ‘vapor’ a subscript in the header for the
second column
Step 6
Click and hold down the mouse button on the top left data item, then drag you
mouse to the lower right hand data item and release the mouse button to
highlight you data table
Step 7
Click on the ‘INSERT’ Tab
Click on the lower right hand corner of the ‘Charts’ box
Click on the ‘All Charts’ Tab
Click on the ‘X Y Scatter’ along the left side of this window
At this point the window should show a little mini-plot of your data with
no lines connecting the points. If this is so, click on the OK button at the
bottom of the window, otherwise, click on the top left chart diagram that
shows points with no connecting lines, then click on the ‘OK’ button
192
At this point you should have a reasonable plot of your data displayed on the spreadsheet.
However the labels for the graph and for the axes are not right. Let’s fix them
Step 8
Axis Titles
Step 9
To correct the Chart Label, simply click somewhere near the Chart label. There
should now be a box made with dashed lines around the label. Click somewhere
in the printing in the label, and add or erase characters until the label is right. It
should read: Data for the Enthalpy of Vaporization of Water.
Now click at some other point in the chart. Do you see a green ‘+’ sign appear
to the right of the graph? Click on that ‘+’ sign to add or remove Chart elements.
Add the chart element ‘Axis Titles’
You should now have the words ‘Axis Title’ added to the left and the
bottom of your two axes. Click on each of the titles and enter the
appropriate title: 1/T in K-1 for the X Axis and Natural log of Vapor
Pressure for the Y axis. You can even add superscripts and subscripts
just as you did earlier.
Adding Gridlines
Step 10
Click on the ‘+’ sign to the right of the graph again
Move your mouse down to the box that says ‘Gridlines’ and click on the arrow to
the right of the box to bring up more options. Click on the boxes for ‘Primary
Minor Horizontal’ and ‘Primary Minor Vertical’. More grid line should have
appeared.
Adjust the Axis Scale
Step 11
Click on the ‘+’ sign appears to the right of the graph. Move your mouse to the
‘Axes’ box and click on the arrow to the right of the box to bring up more
options. Click on the ‘More Options’ box. This will bring up the Format Axis
menu. First click triangle to the right of the ‘Axis Option’ tab near the top of the
menu. Now click on the ‘Horizontal (Value) Axis’ bar. You can now set the
Bounds of this axis. Let’s leave the minimum alone, but change the maximum to
.0031. Now you can go back to ‘Axis Options’ and change the ‘Vertical (Value)
Axis’ maximum to -0.5
Adding a Trendline
Step 12
Click on the ‘+’ sign appears to the right of the graph. Move your mouse to the
‘Trendline’ box. First click on the box, so it will display a trendline, and then
click on the triangle to the right of the box so we can get more options. Click on
the ‘More Options’ box. At the top of the Trendlines Options menu are three
icons. Click on the rightmost icon that looks like three vertical bars. Under the
Trendline Options make sure the button for a ‘Linear’ fit is on, and click on the
box near the bottom for ‘Display Equation on chart’
Note: when you back to the chart you can click on the equation for the line and
move it to a different spot on the graph if you want
193
Note: Another way to access many of the commands given on the previous page is to click
anywhere on the graph first. Notice that two new tabs appear under the green ‘Chart Tools’ box.
Click on the ‘Design’ tab the on then on the far left ‘Add Chart Element’ box. You now have a
different way to get to these commands. Use whichever way you like the best.
To Print Your Spreadsheet
First you have to decide if you want to print just the graph or the graph and the
spreadsheet together.
If you want to print just the graph, click anywhere in the graph, then click on ‘File’, then
click on ‘Print’ on the left hand side of the file menu, and hit the ‘Print’ Box.
If you want to print the spreadsheet data and the graph, first click anywhere on your
spreadsheet. Notice that the spreadsheet has a dashed line through the middle? This
shows where the pages that you will print are separated. The chances are good that your
graph straddles this line, so it will be printed on two different pages. This is not good.
Click on your graph and move it so it is on the same page as your data. You may have
to play with this a bit, because, depending on where you clicked on the graph, it may try
to move a title or an equation. Also, once you have it positioned where you want it,
click somewhere back on the spreadsheet, otherwise the next print command will print
just the graph, and not the entire spreadsheet.
Once you have your page set up, click on File, Click on ‘Print’ on the left and side, and
hit the ‘Print’ box.
194
Names
Section
Report Sheet Excel Graphing Exercise I
A. What is the equation for the trendline for the Enthalpy of Vaporization data?
____________________________
B. What is the slope of this line? __________
Staple a copy of your graph for Exercise I to this report sheet.
195
Exercise II.
Accurately reading the freezing point of a solution from the graphed cooling curve data is
more difficult than reading the freezing point of a pure solvent from its cooling curve graph.
The pure solvent is easy as once it starts to freeze the temperature stays the same, levels off, and
you can read the freezing temperature directly from the y axis. Whereas the cooling curve for a
solution continues to drop as the solution freezes, due to the solute being more concentrated in
the remaining liquid. So to find the “true” freezing point of a solution you need to plot the best
fit line for the data points above where freezing begins, and a second best fit line for data points
after freezing begins. You can use trendlines to fit your lines, but you still need to select the
data points you use for the lines from all the data you gather with your cooling curve. This
means you need to visually assess where the most linear portion of your above freezing data is
and the most linear portion from the below freezing area, as there will be some data in the
middle where the data is curving as it transitions from one slope before freezing starts to a
second slope after freezing starts. Below you will find a table of data from a cooling curve for a
solution of cyclohexane which contains 0.199 g of a solute. Using Excel and to set up a spread
sheet with a data table and start to graph this cooling curve just like the one you did in exercise
I. But once you reach Step 11 , Adding a trendline, we will do a few things differently. Go to
the top of the next page to see detailed instructions.
Data for Freezing Point of Cyclohexane
with 0.199g of a solute added to it
time (sec)
Temp degrees C
0
15.0
15
12.8
30
11.1
45
9.5
60
8.4
75
7.3
90
6.6
105
5.9
120
5.3
135
4.9
150
4.6
165
4.4
180
4.3
195
4.4
210
4.4
225
4.1
255
3.8
270
3.8
285
3.8
300
3.7
315
3.7
330
3.7
345
3.5
360
3.5
196
Once you have reached Step 11 you should have a window containing the initial graph of all the
cooling curve points plotted and the axes labeled and scales appropriately.
Click anywhere on the graph. To the right of the graph are three icons. I’m not sure what the
bottom icon is supposed to be, but click on it. This brings up a new window, click on the
‘Select Data’ box at the bottom right of the window. This brings up the Select Data Source
Window. On the left hand side of this window click on the ‘Add’ button. Under ‘Series Name’
type “Above Freezing”. Next click on the tiny little graph to the right of the ‘Series X values’.
Don’t type anything in, but go to the spreadsheet and right click on your first X value, and , still
holding the mouse button down, drag the mouse to the seventh value (60). See how excel fills
in these the Edit Series box? Now click on the tiny graph to the right of the ‘edit series’ box,
any you will be returned to the main ‘Edit Series’ box. In a similar manner fill in the first 6 Y
values in the ‘Series Y values’ box. Now click the OK button.
Now that you have ‘Above Freezing’ points, in a similar manner add a ‘Below Freezing’ set of
points that includes the values from 225 seconds to 360 seconds.
When you are finished look at your graph. You should now have three different colored points
on your graph. Now click on the graph so you get the ‘+’ sign to the right of the graph. Click
on this and then click on the triangle to the right of the ‘Trendline’ box and click on the ‘More
Options’ dropdown menu. The computer now wants to know which set of data you want to
make a trendline for. Click on ‘Above Freezing’ and the ‘Format Trendline’ Menu will appear.
Click on the rightmost icon (the three bars) near the top of this menu. Make this a ‘Linear’
trendline, and display the equation on the chart. Note that the trendline is just between the
chosen points, and does not extend down. We want it to extend forward so we can see where it
intercepts the trendline from the below freezing points. To do this, find the ‘Forecast’ section
on this menu. Fill in 10.0 period in the ‘Forward’ box. Then click on the graph and see how
this looks. It extended it, but not very far. Now simply click on the trendline, it will bring up
the trendline menu, and adjust the forecast number until you see something you like.
Now add a trendline for the ‘Below Freezing’ data, display the equation on the graph, and
forecast the line backward so you can see where the two lines cross. At this point you may
want to add additional gridlines and change the minimum and maximum on your axes so you
get a high quality graph.
197
Names
Section
Report Sheet Excel Graphing Exercise II
A. Equation for trendline above freezing
Slope of line
Y-intercept
B. Equation for trendline below freezing
Slope of line
Y-intercept
C. Calculate the Freezing Point of this solution two ways
1. By extending the trendlines and the point at which they intersect will give the
freezing point
Freezing Point from graph
2. By mathematically setting the line equations equal to each other solving for x,
then plugging that x value back into one of the line equations to get the freezing
point temperature, the y value. (Show calculation on back of report sheet)
Freezing Point by calculation
Staple a copy of your graph for Exercise II to this report sheet.
198
Using Excel to Graph Data
(Office 2007 version)
Objectives:
<
<
learn to use Excel to graph data
learn how to use the trendline function to plot a line of best fit from a set of
data points
Graphing is an important method of analyzing data gathered from experiments. Graphing can
make it easier to see trends in sets of numbers visually for comparison; or in the case of an
experiment or assay that involves plotting a line of best fit for your data points, it is
indispensable. Today we will use the Chart Wizard in Excel to plot sets of data points, then plot
a line of best fit for the data, and finally have the equation for the line printed out on the plot
(the line equation is written as y = mx + b, where m is the slope of the line and b is the yintercept). Exercise I of this lab will lead you through the process of setting up a graph step by
step. In Exercise II you will set up a graph for a different data set and learn how to designate
data series and apply trendlines to more than one series.
Exercise I.
Step 1
find a computer with Microsoft Excel 2007 installed on it
Step 2
select Excel and open the program
Step 3
under the ‘Page Layout’ Tab select
Size - Letter 8.5" x 11"
Orientation- Portrait
Margins - Custom Margins
Top: 1
Bottom: 1
Left: .75
Right .75
OK
199
Step 4
Enter the following column titles and data into columns A & B on the Excel
spreadsheet. (See data table on next page) You can pull the columns wider by
placing the mouse pointer on the column line and clicking and holding the left
button while you drag the column wider. (Note: Excel will by default use data
in column A for X values and data from column B for Y values in a scatter
plot. And if you have data in columns A, B, C, & D highlighted it will plot
three series using A as the x value for all three and B, C, and D as y values)
Data for Enthalpy of
Vaporization of Water
1/T (K -1)
ln (Pvapor)
.00283
-.799
.00286
-.941
.00288
-1.050
.00291
-1.204
.00293
-1.309
.00296
-1.470
.00298
-1.609
.00301
-1.772
.00304
-1.897
.00307
-2.040
.00310
-2.207
Step 5 left click on the top left cell, while holding the mouse button down move the mouse over
the bottom right cell and release the mouse button. The data table should now be
outlined with a heavy black line and should have a gray background to show it has been
selected.
Step 6
left click on the Insert tab
Step 7
left click on ‘Scatter’ in the charts section
Step 8
left click on the top left chart sub-type (this one shows only data points no lines)
A plot of your data will now appear and the format of the buttons at the top of
the screen will revert to the default Excel buttons. While you have a workable
plot you need to clean up the plot and add important features like titles for the
plot and the axes
200
Step 9
left click anywhere in the plot. Note of the buttons at the top of the screen
change as Excel shifts to ‘Chart Tools’ mode
Step 10
left click on the Layout Tab
left click on the Chart Title icon in the Labels section
left click on the ‘Above Chart’ icon
There will no be a text box on your chart that says ‘Chart Title’
Left click in the box and replace the words ‘Chart Title’ with Enthalpy of
Vaporization for Water
Step 11
left click on the Axis Titles button under the Layout tab
Under Primary Horizontal Axis Title Click ‘Title Below Axis’ and replace the
text ‘Axis Title’ with 1/T in K-1,
Step 12
In a similar manner label the Y axis Natural log of Vapor Pressure.
Step 13
left click on the gridlines icon in the Axes section
Set both Primary Horizontal Gridlines and Primary Vertical Gridlines to Major
& Minor Gridlines (gridlines make it very easy to read values off your graph.)
Step 14
left click on the Legend icon in the Labels section and select ‘None’ to remove
the legend. (In Exercise II you will need to display legends)
Step 15
left click on some white space on your graph, and move it to a good spot on your
Excel spreadsheet so you can print out the data and the graph on a single piece of
paper. For Exercise I drag it down below the data in columns A & B and
place the graphs left edge on the left margin of the sheet.
Step 16
left click and hold on the lower right corner (you should get a double ended
diagonal arrow) and drag the corner out and down to fill up the page with the
graph. To preview the page select the sheet by left clicking on the shaded cell to
the left of the A column label then select File and print preview. (If you don’t
select the sheet the preview will only show you the graph, not your data table or
header.) If you are happy with the placement of the graph on the sheet, close
window to return to chart.
Now your graph is pretty well set up and all that is left is fine tuning to make it easier to read
values from
Finishing up the graph
201
To set X and Y axis scales
place mouse pointer on the Y axis of graph right click and select format axis,
for Exercise I find the maximum row, click on the ‘fixed’ button and set the
maximum to -.5.
Next adjust the X axis. The scale here is OK, but let’s move the numbers above
the grid. Place the mouse on the X axis of the graph and right click and select
format axis. In the axis options box change axis labels: to ‘High’. Depending on
the size of your plot you may want to change the format of the number (under
Format Axis - Number option) or the size of the font (Window that comes up
when you first right click on the axis)
To adjust placing of graph title or axis labels left click on the area of interest
<
to move area left click and hold while you drag area to where you want it.
<
to correct spelling or add text left click in the text to be changed to drop the text
cursor
<
to change font size right click to get into the format window
If you need to have a line of best fit plotted
place mouse pointer on one of the data points on graph
right click
left click on Add trendline
In the Trendline Options Window
Click on the ‘Linear’ button
select display equation on chart box
left click on CLOSE (This will plot the best fit line on your graph and
print the equation for the line on the graph. The equation for a line is
y=mx+b, where y and x are the axis values, m is the slope and b is the yintercept.)
To Print Your Spreadsheet
<
select the spreadsheet by left clicking in the shaded cell to the left of the
cell labeled A (if you don’t select the whole sheet it will only print your
chart)
<
left click on the Windows Office icon in the upper left hand corner
<
move your mouse over the print icon, and then click on Print Preview
<
if preview is what you want left click on print
202
Names
Section
Report Sheet Excel Graphing Exercise I
A. What is the equation for the trendline for the Enthalpy of Vaporization data?
____________________________
B. What is the slope of this line? __________
Staple a copy of your graph for Exercise I to this report sheet.
203
Exercise II.
Accurately reading the freezing point of a solution from the graphed cooling curve data is
more difficult than reading the freezing point of a pure solvent from its cooling curve graph.
The pure solvent is easy as once it starts to freeze the temperature stays the same, levels off, and
you can read the freezing temperature directly from the y axis. Whereas the cooling curve for a
solution continues to drop as the solution freezes, due to the solute being more concentrated in
the remaining liquid. So to find the “true” freezing point of a solution you need to plot the best
fit line for the data points above where freezing begins, and a second best fit line for data points
after freezing begins. You can use trendlines to fit your lines, but you still need to select the
data points you use for the lines from all the data you gather with your cooling curve. This
means you need to visually assess where the most linear portion of your above freezing data is
and the most linear portion from the below freezing area, as there will be some data in the
middle where the data is curving as it transitions from one slope before freezing starts to a
second slope after freezing starts. Below you will find a table of data from a cooling curve for a
solution of cyclohexane which contains 0.199 g of a solute. Using Excel and chart wizard set up
a spread sheet with a data table and start to graph this cooling curve just like the one you did in
exercise I. But once you reach Step 8 , the initial graph of the data, we will do a few things
differently. Go to the top of the next page to see detailed instructions.
Data for Freezing Point of Cyclohexane
with 0.199g of a solute added to it
time (sec)
Temp degrees C
0
15.0
15
12.8
30
11.1
45
9.5
60
8.4
75
7.3
90
6.6
105
5.9
120
5.3
135
4.9
150
4.6
165
4.4
180
4.3
195
4.4
210
4.4
225
4.1
255
3.8
270
3.8
285
3.8
300
3.7
315
3.7
330
3.7
345
3.5
360
3.5
204
Once you have reached Step 8 you should have a window containing the initial graph of all the
cooling curve points plotted as, series 1. What you will need to do now is visually decide which
data points to use for the two lines you need to calculate freezing point of a solution. It looks
like the first six data points will work for the above freezing line and the last nine for the below
freezing line. To enter a series of data for the above freezing line:
click on the Select Data icon in the Data section of the Design Tab:
click the Edit button on the left half of the Select Data Source Window
Change the Series name to Full Cooling Curve
Hit the OK button
click the Add button on the left half of the Select Data Source Window
Click in the Series Name and make the name of this series “Above Freezing”
Click in the Series X value box - then click and drag the mouse over the first six
X values in the data table
Click in the Series V value box - delete the “={1}” then click and drag the mouse
over the first six Y values in the data table
Hit the OK button
click the Add button on the left half of the Select Data Source Window
Click in the Series Name and make the name of this series “Below Freezing”
Click in the Series X value box - then click & drag on the last nine X values in
the data table
Click in the Series V value box - delete the “={1}” then click & drag on the last
Y values in the data table
Hit the OK button
Click on the OK button - The Select Data Source window will disappear you your graph
should appear with your three different regions defined in different symbols and colors
Now that you have your data series all entered you can continue on with Step 9 and
finish the graph.
Once the graph is dropped onto you spreadsheet then you can right click on a data series
point to put in a trendline and line equation, for the above freezing series and another
one for the below freezing series. And at the same time you can change colors of the
data points if you need to.
Remember to set up the scale and gridlines so that you can read Temperature to tenths of a
degree on the y-axis. (Minor gridlines at 0.5 will be fine for this.)
205
Names
Section
Report Sheet Excel Graphing Exercise II
A. Equation for trendline above freezing
Slope of line
Y-intercept
B. Equation for trendline below freezing
Slope of line
Y-intercept
C. Calculate the Freezing Point of this solution two ways
1. By extending the trendlines and the point at which they intersect will give the
freezing point
Freezing Point from graph
2. By mathematically setting the line equations equal to each other solving for x,
then plugging that x value back into one of the line equations to get the freezing
point temperature, the y value. (Show calculation on back of report sheet)
Freezing Point by calculation
Staple a copy of your graph for Exercise II to this report sheet.
206
Constants
Avogadro's number
Boltzmann constant
Faraday constant
Gas constant
N
k
F
R
Plank's Constant
Speed of light
h
c
6.022×1023 mol-1
1.38066×l0-23J/K
96,485 C/mol
8.31451 J/K@mol
0.08206 L@atm/K@mol
6.62608×l0-34 J@s
2.99792458×l08m/s
Conversion Factors
Length
1 meter = 1.0936 yards
1 centimeter = 0.3937 inch
1 inch = 2.54 centimeters (exactly)
1 kilometer = .62137 mile
1 mile = 5280 feet
Volume
1 liter = 1.0567 quarts
1 gallon = 4 quarts
1 quart = 32 fluid ounces
Energy
1 joule = 1 kg@m2/s2
= .23901 calorie
1L@atm= 101.3 Joule
Mass
1 kilogram = 2.2046 pounds
1 pound = 16 ounces
1 ton = 2000 pounds
1 metric ton = 1000 kg
Pressure
1 pascal = 1 N/m2
= 1 kg/(m@s2)
l atm = 101.325 kPa
= 760 torr
= 14.7 1bs/in2
1 bar = 105 Pa
207