CHEM 112 Lab Manual Spring 2024
Content
Page(s)
Lab Policies
Laboratory Safety Rules
Pregnancy Notice
Equipment Check-in List
Review of Proper Equipment Use
Academic Honesty & Lab Grades
Lab Report Guidelines
Making Graphs in Excel
Week 1: 1/16 – 1/18
NO LABS FIRST WEEK OF CLASS
Week 2: 1/22 – 1/25
Orientation & Check-in
Week 3: 1/29 – 2/1
Experiment 1: Organic Structures & Nomenclature
Week 4: 2/5 – 2/8
Experiment 2: Gas Chromatography, Intermolecular
Forces & Boiling Point
Week 5: 2/12 – 2/15
Experiment 3: Net Ionic Equations
Week 6: 2/19 – 2/22
Experiment 4: Rate Law for Bleaching Crystal Violet
Week 7: 2/26 – 2/29
Experiment 5: Le Chatelier’s Principle
Week 8: 3/4 – 3/7
Experiment 6: Luminol Lab Practical
Week 9: 3/11 – 3/14
NO LABS – Spring Break
Week 10: 3/18 – 3/21
Experiment 7: Spectrophotometric Determination of
Cu in a Penny
Week 11: 3/25 – 3/28
Experiment 8: Synthesis of an Oxalate Hydrate
Week 12: 4/1 – 4/4
Experiment 9: Percent Oxalate Determination by
Titration
Week 13: 4/8 – 4/11
Experiment 10: Bleach Titration Lab Practical
Week 14: 4/15 – 4/18
Experiment 11: Redox Reactions & Equations
Week 15: 4/22 – 4/25
LAB MAKE-UP WEEK
Week 16: 4/29 – 5/2
Check-out & Lab Clean-up (mandatory)
2
3-5
6
7
8-9
10
11-12
13-15
NA
1-15
16-22
23-27
28-35
36-41
42-50
51-52
NA
53-60
61-62
63
64-66
67-71
NA
72
If you have an excused absence from a lab, it is your responsibility to initiate contact
with your instructor to get an OK for making up the lab with another section OR Lab
Make-Up Week, then you must contact the instructor of the other section to get an OK
from them. Contact concerning lab absences should be made a week in advance for
predictable absences and within 24 hours of your lab period for unpredictable
absences.
Spring 2024 CHEM 112L Lab Sections
Section
001
002
003
004
005
Day
Monday
Monday
Wednesday
Wednesday
Thursday
Time
1-2:50
3-4:50
1-2:50
3-4:50
2-3:50
pm
pm
pm
pm
pm
Room
444
444
444
444
444
Instructor
Z. Lee
Z. Lee
A. Macintosh
D. Fulmer
D. Fulmer
1
Instructor Email
z.lee@moreheadstate.edu
z.lee@moreheadstate.edu
a.macintosh@moreheadstate.edu
d.fulmer@moreheadstate.edu
d.fulmer@moreheadstate.edu
Lab Policies
1. You must attend each scheduled laboratory session. If you have a documented, excused absence,
bring the documentation to your instructor before missing lab if possible and as soon as possible
otherwise. Don’t wait until the next lab. Missing 3 lab sessions or not turning in lab reports
equivalent to 3 lab sessions will result in an automatic E in Chemistry 112. The smooth, safe
operation of lab depends on your hearing prelab instructions and safety cautions. If you are late,
your instructor may assess a grade penalty.
2. Your lab grade will be based on the quality of your experimental work based on your written lab
report, on the quality of your report in terms of communication and/or a quiz. No lab reports will be
accepted for experiments which you did not attend and do conscientious work. Sometimes you may
be given a quiz instead of submitting a report (this will be announced ahead of time).
3. All lab reports are due at the beginning of the lab period immediately following the day an experiment
is completed; they must be done using a word processor unless otherwise stated . All lab reports
are expected to be original and done by individuals, not pairs; this includes any tables, graphs,
etc. If your instructor tells you that you may write a report with your partner, then both partners are
responsible for all of the report. Having one partner unavailable or data unavailable is not an excuse
for no report.
4. When working in the lab, you must follow all the safety rules. When you leave the lab, your work
areas must be cleaned and all equipment returned to its original place. The last 4 persons leaving the
lab are responsible for its overall cleanliness and orderliness. Failure to follow safety rules or
maintain cleanliness and order will result in a 50% penalty for that lab. If you violate safety rules,
your instructor may eject you from lab and give you a zero for the day.
5. Lab notebooks. You are to keep a bound lab notebook with data in chronological order. Record data
in ink and strike through any mistakes with a single line–to not obliterate the data.
2
Name:_______________________________ Course:___________
(print)
Date:___________
LABORATORY SAFETY RULES
1)
Responsible behavior is essential. Think about the consequences of your actions or inactions.
Make note of the safety precautions written in the lab text and told to you by the laboratory
instructor. If you are unsure about a procedure, ASK!
2)
WEAR APPROVED EYE PROTECTION AND AN APRON AT ALL TIMES IN THE LAB.
Goggles and safety glasses are available in the bookstore. If you wear glasses, be sure your
eye protection fits over your glasses. If you must wear contacts, then you must wear goggles.
Your shoes and clothing should also provide protection, i.e. wear closed-toe shoes whose
"uppers" repel spills. Wear clothing that covers you from the ankles to your throat, with
the equivalent of a crew-neck t-shirt on top. Failure to do so will result in expulsion from
lab. Confine long hair and loose clothing. Leggings and yoga pants do not count as pants. They
fit too closely to repel spills. So no bare midriffs sleeveless shirts, shorts, short skirts,
high-heeled or slick-soled shoes.
3)
In case of fire, accident, or chemical spill, notify the instructor at once.
a) Wet towels can be used to smother small fires.
b) In case of a chemical spill on you body or clothing, wash the affected area with large
quantities of running water. Be sure to first remove the clothing that has been wetted by
chemicals to prevent further reaction with the skin.
c) If you should get a chemical in your eye, wash with flowing water from the eye wash for at
least 15 minutes. Get medical attention immediately.
d) Report all injuries immediately to your instructor.
4)
Note the locations of the safety equipment listed below. Make sure you know how to use them.
a) Eye Washes b) Safety Shower
c) Fire Extinguisher
d) Blanket
e) Telephone
f) Sinks
g) Hoods
h) Exits
5)
EMERGENCY HELP IS AVAILABLE BY DIALING 911. Telephones are available in the C
hallway between 445 & 446 and in the middle (B) hall outside 423.
6)
Do not eat, drink, or chew on anything in the laboratory. This includes food, gum, tobacco,
pens, and other chemicals. Also, you should not handle many lab chemicals. They may absorb
through the skin or be ingested if you lick your fingers. Wash thoroughly before you leave the
lab, even if you haven't "touched" anything.
7)
Avoid breathing fumes of any kind.
a) To test the smell of a vapor, "waft" the vapor toward you. Test only those vapors specified.
b) Work in a hood if there is a possibility that noxious or poisonous vapors may be produced. If
unexpected vapors are produced during an experiment, move the experiment to a hood
immediately.
8)
Keep your work area clean at all times. Clean up spills and broken glass immediately, whether
at your bench or in the common area. Clean your work space, including wet-wiping the bench
top and putting away all chemicals and equipment, at the end of the period.
9)
Perform no unauthorized experiments. This includes using more reagent than instructed.
3
10)
Never work alone in the laboratory.
11)
Always pour acids into water when mixing. Otherwise, the acid can violently spatter. "Acid into
water is the way that you oughter."
12)
Do not force a rubber stopper onto glass tubing or thermometers. Use a split-hole stopper if
possible. If not, lubricate the tubing and the stopper with glycerol or water; use paper or cloth
toweling to protect your hand; and grasp the glass close to the stopper.
13)
Beware of hot iron rings and hot glass. They "look" cool long before they can be handled.
14)
Dispose of small quantities of excess non-toxic liquid reagents by flushing them down the sink
with copious quantities of water. Dispose of solids and toxic liquids in the waste containers
provided. Never use your eyedroppers to remove liquids from a reagent bottle. Always pour
reagents into your own clean containers. Never return reagents to the bottle.
15)
Spatters are common in general chemistry laboratories. Test tubes containing reacting mixtures
or those being heated should never be pointed at anyone.
16)
Carefully read the experiment before coming to lab. Review techniques to be used and safety
precautions to be taken.
17)
Finally, and most importantly, THINK about what you're doing. If you doubt that what you're
doing is safe, don't do it. Check with your instructor. Ask questions.
Minimization of exposure of all students to the chemicals we work with in lab is an important
goal for both the instructor and the students. The experiments we pick, chemicals we choose
to use, and safety procedures we emphasize are all a part of attaining that goal. However, be
aware that the effects of many chemicals are different for different people. Certainly, we all
know about allergies, and have heard about how different medicines can interact when taken
together and cause disastrous results. Analogously, certain medical conditions such as
asthma, allergies, etc. can be worsened by exposure to certain chemicals. Please inform your
instructor about any allergies or medical conditions of which they should be aware. Some
chemicals which are not generally considered toxins are reproductive toxins. They may affect
the fetus directly or affect the ability of the parent (fathers, too) to produce healthy offspring.
Also realize that not all chemicals have been tested to see if they are reproductive toxins.
Anyone who is sexually active needs to be diligent in their efforts to minimize their chemical
exposure in the lab. If you are or become pregnant, you need to notify your instructor as soon
as possible. Notification forms are available in your manual or from your instructor.
...........................................................................................................................................................
I have read and understand the Laboratory Safety Rules and have retained a copy for my
reference.
________________________________
(signature)
________________
(course & section)
__________
(date)
Please list any allergies or medical conditions (including pregnancy) of which the instructor should be
aware. Please indicate if you wear contacts.
___________________________________________________________________________
Revised: RKC 8/17
4
Lab Attire Rules
NOT APPROVED:
•
•
•
•
•
•
•
No leggings, no jeggings, no skinny jeans (or similar skin tight pant by a different name), no
shorts, no capris, no ripped/holey pants, no yoga pants. Legs will be totally covered by a loose
fitting, durable material. Just a lab apron or coat is insufficient.
No tank tops, no sleeveless tops, no cropped shirts, no plunging necklines, no spaghetti straps,
or ripped shirts.
No Sandals, no open toe, open back, or open weave shoes, shoes with holes in the top or sides,
no Birkenstocks, Mary Janes, or Crocs
No hats or caps
No head phones or earbuds
No “forehead protectors” (your eye protection comfortably positioned on your forehead).
The instructor reserves the right to determine inappropriate clothing at any time that is not listed
here. If you are curious about a particular piece of clothing, consult your instructor a minimum
of 24 hours before lab.
APPROVED MUST HAVES:
•
•
•
•
•
You must wear a top that covers you upper torso and shoulders equivalent to short sleeves (a
good example is a crew neck t-shirt, that fits properly).
You must wear pants that covers your lower torso from your waist all the way to the top of your
shoes (a good example is a regular pair of jeans or kaki’s that fit properly)
You must wear shoes that cover you whole foot and fit correctly, (your no mesh tennis shoes)
You must wear a plastic/vinyl apron or lab coat
You must wear colorless safety goggles or safety glasses. Safety glasses must be approved by
your instructor. They must have side and top shields. VisorGogs are a nice option for some as
they breath better than traditional goggles.. The goggles will indicate some sort of code relating
to an ANSI Z87.1 Standard (ANSI Z87.1-1989, ANSI Z87.1-2003, ANSI/ISEA Z87.1 D3-2010,
ANSI/ISEA Z87.1-2015 or Z87+ D3).
If you do not abide by these rules at any time during your lab period you will be
excused from lab with zero credit earned. Time is not available to run back to
residence for appropriate attire. When in doubt, simply carry scrubs or pants to
wear over shorts while in lab.
5
Notification of Pregnancy
I hereby notify my lab instructor that I am pregnant. I give permission to the laboratory
supervisor to send to my doctor a list of the chemicals I will be using this semester. I also authorize
my doctor to discuss with my instructor or the laboratory supervisor any hazards associated with
exposure to these chemicals. I also authorize my doctor to respond in writing to the lab supervisor
with recommendations for extra precautions they recommend, including exclusion of specific
chemicals from the experiments I perform.
Student's Name ________________________________________
(please print)
Student's Signature _____________________________________
Student's Social Security # _________________
Student's Date of Birth ____________________
Doctor's Name _________________________________________
Doctor's Address _______________________________________
_______________________________________
_______________________________________
Doctor's Phone # ________________________
Course _______________
Lab Instructor _________________________________
Date _________________
6
Equipment List for Chemistry 112
Name(s)
Drawer # __________ Date _______________
Lab section or lab day/time ____________________
In week one, check your desk to see that you have the following items, clean and unbroken.
Your instructor will tell you where to place extra items. Replace any broken or missing
equipment. Then keep this page for the last day of lab.
1- 20mL beaker
2- 100 mL beaker
2- 250 mL beakers
2- 250 mL Erlenmeyer flask
2- scoopula
1-test tube clamp
10 test tubes, 13 mm x 100 mm
2- Pasteur pipet (looks like a long medicine dropper),
with bulb
2- 50 mL beakers
2- 150 mL beaker
1- 400 mL beaker
2- funnel
1- triangle
1 test tube rack
1- pair crucible tongs
2- glass stirring rods
1- 500 mL plastic bottle for Crystal
violet (stained is OK)
1- 100 mL graduated cylinder
1- 10 mL graduated cylinder
2- 250 +/or 125 mL plastic bottles
Check the large locker by your drawer for the following:
one ring stand
one ring
one utility clamp
one wire gauze
one bunsen burner and tubing
Keep this page for the last day of lab!
******************************************************
Check out:
Student’s signature(s):
Date
Instructor’s signature
7
Review of Proper Equipment Use
Many labs direct you to use YouTube videos and to review the proper use of various pieces of
equipment before lab. Below are links to videos associated with some important glassware and questions
you should be able to answer about each.
Volumetric Pipet
Use of a pipet: https://www.youtube.com/watch?v=TffTiRw8cQY
1. Be able to describe the procedure used to rinse the pipet properly.
2. Where should the tip of the pipet compared to the liquid surface while liquid is being drawn into
the pipet? While liquid is being let down to the line on the pipet stem?
3. Which of the following shows the proper adjustment of liquid level in a pipet so it will deliver
its stated volume?
4. What do you do about any droplets or partial droplets of liquid on the outside of the pipet?
5. How much liquid should be left in the tip of the pipet after it gravity drains?
6. Where should you look to find the precision of the pipet (so you know how many decimal places
to record in the volume)?
7. What should be done with the pipet before it is returned to the pipet storage container? (rinse it!)
Volumetric Flask, Making and Diluting Solutions
Use of a volumetric flask to make a solution: https://www.youtube.com/watch?v=Btp1N5Z2L74
• Video notes:
o You can also dissolve a solid in a small beaker and then add it to the flask (but then
rinse the beaker into the flask).
o He has already rinsed the pipet well.
o Invert the flask more times than he did.
1. George needs to make 250.0 mL of an aqueous solution. He weighs the appropriate amount of
solute needed to make 250.0 mL of the solution, adds 250.0 mL of water to it, stirs to dissolve
the solid and mixes it well. What did he do wrong?
2. What should you do with the volumetric flask before using it?
3. Which of the following shows the proper liquid level in a full volumetric flask that is being used
for quantitative purposes?
4. At what points should you mix the solute and solvent by shaking the flask well? Why should you
do this both before the liquid is filled to the line and then again after? Why should the flask be
inverted during some of this mixing?
5. What should you do with the flask when you finish with it?
8
Buret
Use of a buret: https://www.youtube.com/watch?v=9DkB82xLvNE
• Video notes:
o It says drain to 0 mL or below. Don’t waste your time getting to exactly 0.00
o You must also rinse the buret well when you are finished with it, including the tip.
o You do not have to add partial drops in this CHEM 111.
o The latter part of this video is about doing a titration (which is where burets are
commonly used.)
Carrying out a titration: https://www.youtube.com/watch?v=9DkB82xLvNE
1. Be able to describe the procedure used to rinse the buret properly.
2. What should be done about any air in the tip of the buret (below the stopcock)? When should this
be done?
3. If a buret is calibrated with lines every 0.1 mL, how many decimal places should be recorded in
the volume reading?
4. Why is it desirable to (a) not use a very small volume of titrant during a titration and (b) not use
a large enough volume to need to refill the buret during one titration?
5. Sometimes a student misjudges and has to refill the buret during a titration. How do you decide
if it is time to refill the buret? What do you do in the next trial to avoid having to refill the buret
in it?
6. Your lab manual will often suggest (or assume that you know by now) first doing a small, sloppy,
trial titration and scaling the volumes appropriately when you start the titration in which data is
being carefully collected. Why not just be careful from the beginning and not do the sloppy
titration?
7. What should you do with the buret when you are finished with it?
9
Academic Honesty & Lab Grades
A Statement on Academic Dishonesty
(written by D. Herron, former Chair, Dept. of Physical Sciences)
“If science is to produce new knowledge about the world in which we live, it is essential that
scientists adhere to the highest possible ethical standards. If scientists are forgiven for reporting false
data, disregarding “bad” experiments, or making up d ata rather than performing experiments, it is
impossible to trust any of their results. This goes deeper than abstract principles of justice and honesty;
the whole scientific enterprise depends on intellectual honesty, and part of what you must learn is
how important absolute honesty is in science. Consequently, any incident of a) falsified data, b) “dry
lab” reports, or c) plagiarism of other peoples results will be treated as a serious breach of ethics and
handled according. The penalty for such behavior will likely result in a minimum penalty of zero for
the experiment at hand, and more serious penalties such as a directed grade of E for the course or a
recommendation for expulsion from the university may be considered.”
In Chemistry 112 labs, you are to write your own lab reports. This includes making your own
graphs and tables as well as writing text. You are also responsible for keeping anyone else from
using your graphs, tables and text, so don’t lend them to someone “just to see”. You may talk to your
lab partner about your data, etc., but you should not write down things while you are doing it. If you
write as you talk, your lab reports will sound too much alike and then your instructor is likely to accuse
you of cheating. So, “I worked with him/her” is not an acceptable reason for reports being too much
alike. The penalty for copying from someone or someone else copying from you will range from a zero
on the report to an E in Chemistry 112.
The laboratory grade counts for 15 % of the Chem grade. However you must pass both laboratory and
lecture with the minimum grade to pass. This means that not passing either laboratory or lecture will
result in a failing grade for the entire class. If you miss a lab with an excused absence you are expected
to make the lab up or you will receive a zero. If you miss three labs without them making up or if you
do not turn in the reports for the equivalent of three days of lab (which may or not be three lab reports),
then you will receive an automatic E in Chemistry 112.
I have read and understand the above rules; I agree to abide by them and I understand the penalties.
_______________________
Signature
_________________
Date
Sign this sheet and turn it in to your instructor.
10
Lab Report Guidelines
Generic Report Format
The format below is general. Specific experiments may dictate some variations in report format, but
all should follow this general form. Do not use a cover sheet.
Top of first page:
• Your name, your partner’s name if you worked with someone (collecting data)
• Chem 112
• Lab day, time
• Title of the experiment (title of what you did, not a generic title). For example, Analysis of
Draino, not Analysis of Household Cleaner.
Objective:
This is one sentence stating exactly what specific task you did in this laboratory exercise.
Procedure and Observations:
A detailed description of the work you did. This includes all observations, generally mixed in with
the procedure. If you are doing quantitative tasks and follow exactly the procedure in this manual, then
just state that you did so; but you must still include observations and enough detail to tell when you saw
various observations. Do not skimp on observations. If a change in the procedure is given, you must
tell this. Give any procedure which you devise.
Results:
Data Tables: For data which differ from trial to trial but are collected using the same general
procedure, use standard word processor tables or (if you need to do calculations on tabulated data)
spreadsheet tables. The tables should be as compact as possible while being easy to understand. Do
not break a table over two pages. All data must be clearly labeled. If any data is clearly incorrect
(if there is reference book data available or if there are duplicate measurements), then indicate this
in some way and then do not use it in calculations.
Sample Calculations: If a calculation is based on a chemical reaction equation, write it down.
Summarize the calculation with a “word equation” first, then show one calculation with numbers
substituted in. Show one of each type of calculation this way; results of repeated calculations of the
same type should simply be placed in a table and then the average and standard deviation should be
reported for that type of calculation.
Discussion and Conclusions:
This section states what you found and how well you achieved your objective. Your conclusion
should be directly related to the stated objective in your report. Leave out the filler (BS). You should
support this only with specific data and observations which are included above in the report. Of
course, you should also report specific experimental error(s), how they affected your results, and how
the results might have been improved, but none of the “I might have misweighed” sort of errors
which you have no evidence about. If possible, find the “correct” answer in a handbook or in a
reliable on-line source (properly cite where you find it) and calculate a percent error.
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Common Things to Watch Out for with Lab Reports
1. The writer should read the report after it is written and revise it as needed. The order of the report
may or may not be the order in which you did things in lab.
2. Don’t write a week-by-week report (i.e. Don’t write “Week 1”, etc.) Write a report for one project
unless the structure of the experiment is so you are really doing many unrelated activities.
3. Don’t write instructions of how to do the experiment. Write a report of what you did. The
admonishment “Someone should be able to do the experiment from your report” means that you
should have sufficient detail in your report, not that it should be a set of instructions.
4. Report repetitive data in tables. Describe a procedure done over and over only once and then list
the samples for which it was used.
5. Pay attention to significant figures!
6. Organize your report so it is easy for the reader to follow what is done. For example, if step B of a
procedure can’t be done until step A has been done, report A first so your reader doesn’t say to
themselves “What about A?” and then later “Oh, there’s A!” A second example: data used in
more than one table should never be simply referenced as “sample 1”, etc. in a manner that makes
a reader have to look back-and-forth between the two tables. Sections of the report should be
similarly organized—don’t put the conclusions before the data and analysis, for example. This is
reader-unfriendly and shows poor communication skills.
7. Don’t connect the dots on the graph.
8. Look up appropriate handbook data to check your experimental data. Calculate averages from
final answers of individual trials, not from raw data. (For example, calculate the molarity of each
trial solution, then calculate an average molarity rather than averaging the milliliters used and
calculating only one molarity.) If you have more than one answer, calculate an average and a
standard deviation.
9. We will assume good lab technique for things like washing glassware (so don’t bother to tell us
that you did it).
10. Run the spell checker on your report (though don’t believe everything the spell checkers suggest);
in particular, the term absorbance is preferred to absorbancey and HCl is not Hcl. Use subscripts
and superscripts where appropriate.
12
Making Graphs in Excel
Directions for PC Computers
Different computers have different defaults of what menus show, etc.. You might need to figure out
how to get to the commands. While Microsoft Office products use the term chart, we may often use the
term graph, which is more specific. I sometimes alternate below, depending on whether I am quoting
what an Excel menu says or not.
Getting the “raw” graph of one set of data:
• Type your independent value (x) data in one column and your y-data in the column immediately
to its right. If it is already typed but not side-by-side or in the wrong order, it is probably easiest
to just cut-and-paste it into the correct order.
• Select all of the data by dragging the cursor over it.
• Choose Insert. From the icons that appear, choose Scatter and then from the popup menu, choose
the icon that has only dots—no lines or curves. The graph should now appear, but you must edit
it before turning it in!
Required Editing of the “raw” graph:
• Click in the middle of the graph to select it.
• Make sure the background is white. If it is not, put the cursor somewhere on the graph but away
from points and lines. Right click and choose format chart area. No Fill should be selected.
• Make sure the graph has a title either above the graph or below the graph. The title should tell
what is being graphed. An example would be Graph 1: Volume of Sample vs. Mass of Sample.
No units are given in a title. With the graph selected, choose Layout Chart Title>Above Chart.
Now you can select the generic title that has been inserted and edit it.
• Axes should be labeled, and the label should include a unit unless the value being graphed has
no unit (which is rare). With the chart area selected, choose Layout and then Axis Titles>Primary
Horizontal Axis Title>Title Below Axis. Now choose Layout>Axis Titles>Primary Vertical
Axis and either Vertical title or Horizontal Title (you might try both and see which looks better).
Now the axes should have generic titles “Axis Title” shown on the graph. Click on these to edit
them.
• If there is only one type of data on the graph, you should remove the legend. Choose
Layout>Legend>None.
• Adding a Line of Best Fit and R 2
o If the points on the graph appear to fall pretty much along a straight line, then add a bestfit line: select Chart Tools>Layout>Trendline>Linear Trendline. Also, after this select
Chart Tools>Layout>Trendline>More Trendline Options. Select Display Equation on
Chart and Display R-squared value on chart.
Editing that Might Be Needed
• Zooming in on one part of the data. If Excel defaulted to showing the intersection of the axes
at (0,0) and your data is considerably away from the 0 point, you should “zoom in” so the data
points are distributed over most of the page. Occasionally you might want to show not only a
general graph of the data, but you may want to have a second graph that shows only a section of
the data.
o To display only a section of the data based on the x values, choose Layout>Axes>
Primary Horizontal Axis>More Primary Horizontal Axis Options. In the popup window
that appears, for Minimum choose Fixed and then enter the value of the leftmost
13
•
•
gridline/axis tic mark that you want to appear on the x-axis of the zoomed graph. For
Maximum, choose fixed and enter and the value of the rightmost gridline/axis tic mark
that you want to appear. You may repeat this process with the y-axis if you desire by
choosing Primary Vertical Axis in step one.
Showing More Tic Marks on the Axes. Choose Layout>Axes> Primary Horizontal
Axis>More Primary Horizontal Axis Options and Uner Major unit chose Fixed and then enter
the distance (in values) that you want between the tic marks. If you want to display some tic
marks without numbers, you may fill in the minor tic mark entry as well.
Changing the number of decimal places shown in the numerical labels on the axes.
Layout>Axes> Primary Horizontal Axis>More Primary Horizontal Axis Options. On the left
side, choose Number. Then choose number in the right list. Enter the number of ecimal places
in the box at the top right.
o There are many other options available under Layout>Axes> Primary Horizontal
Axis>More Primary Horizontal Axis Options.
Directions for Macintosh Computers
Different computers have different defaults of what menus show, etc.. You might need to figure
out how to get to the commands. While Microsoft Office products use the term chart, we may often
use the term graph, which is more specific. I sometimes alternate below, depending on whether I am
quoting what an Excel menu says or not.
Getting the “raw” graph of one set of data:
• Type your independent value (x) data in one column and your y-data in the column immediately
to its right. If it is already typed but not side-by-side or in the wrong order, it is probably easiest
to just cut-and-paste it into the correct order.
• Select all of the data by dragging the cursor over it.
• Click on Charts > Scatter>Marked Scatter
• Click in the middle of the graph, away from any numbers or symbols.
• Chart Layout>Chart Title (choose above or below and then type in a title specific to the graph)
• Chart Layout>Axis Titles>Horizontal Axis Title (then choose one and type in the x-axis title,
including units)
• Chart Layout>Axis Titles>Vertical Axis Title (then choose one and type in the x-axis title,
including units)
• If there is only one set of data, delete the Legend by clicking in it and selecting return
• If the data is linear, add a best fit line by Chart Layout>Trendlines>Linear. Repeat Chart
Layout>Trendlines and choose Trendline Options>Options, then click on Display Equation and
Display R2 value.
For each axis:
• If the first point is very far from the number where the axis starts (excess white space), examine
where the first point is numerically. Then double click right on the axis. In the box that comes
up, choose scale and type in a whole number slightly less that that of the first point’s value on
that axis.
Double-click one of the points and unclick Shadow if it is on.
14
Putting More Than One Set of Data on a Graph Excel
This will most commonly occur with the two sets of data that have the same x values and different
y values, this is assumed in the instructions below.
• Enter all the x values in one column, and each set of y values in its own column. So, for
example, x values in column A, first y set in column B, second y set in column C, etc.
• Now select all of the data by dragging over it and then do the steps as you would if you were
graphing only one set of data and had the data selected. This should result in a graph where
each set of data has a separate set of symbols and Excel know to treat each separately (when a
line is added, for instance)
• Generally if there is more than one set of data on the graph, you should keep the legend and
label it appropriately.
o Changing the labels in the Legend: choose Chart Tools>Design>Select Data. On the
popup window, under Legend Entries select the name of the series that is being given
in the legend (the one you want to re-label). Then choose Edit and under Series Name
enter the name you want to appear on the legend. Click OK until you exit.
15
Experiment 1: Organic Structures & Nomenclature
Introduction
This exercise is designed to introduce you to the naming and structure of organic molecules. As
you work through this exercise keep the following basic fact in mind:
•
IN SIMPLE ORGANIC COMPOUNDS, CARBON SHARES FOUR AND ONLY FOUR
PAIRS OF ELECTRONS!
Atoms are bonded in organic compounds as follows:
• C shares four pairs and has no lone pairs. ALWAYS.
• H shares one pair of electrons with any atom. ALWAYS.
• O shares two pairs of electrons, and has two unshared pairs. ALMOST ALWAYS.
• N shares three pairs of electrons, and has one unshared pair. ALMOST ALWAYS.
• F, Cl, Br, and I share one pair of electrons and have three unshared pairs. ALMOST
ALWAYS.
NOTE: It is incorrect to refer to a pair of electrons that is in a double or a triple bond as “a bond”;
that is two pairs of electrons in C=C is refered to as one bond, not two.
Organic compounds are represented by chemical formulas. A molecular formula indicates the kinds
of atoms and the number of atoms (e.g., C27 H46O), whereas a structural formula simply describes which
atoms are bonded to which other atoms and how the atoms are bonded. The structural formulas are often
written as some sort of condensed formulas (also called condensed structures). For example, a
compound with the molecular formula C10 H22 may be written in the following ways:
condensed structural formulas
H H H H H H H H H H
H C
C
C
C
C
C
C
C
C
C H
H H H H H H H H H H
or
CH3(CH2)8CH3
CH3CH2CH2CH2CH2CH2CH2CH2CH2CH3
or
structural formula
line (stick) structure
Another possibility for C10 H22 is:
CH3
CH3
CH3CHCH2CH2CH2CCH3
CH3
Notice that these two C10 H 22 compounds both have the same molecular formula, but the atoms are
bonded differently. The first is a straight chain compound and the second is a branched chain compound.
These compounds are called skeletal isomers of each other. There are 73 other isomers of C10 H22 !
16
3-D drawings use a plain line for a pair of electrons in the plane of the paper, a solid wedge with its
narrowest end on the central atom to indicate a bond pointing towards the viewer, and a dashed wedge
to indicate a bond pointing away from the viewer:
C
H
solid wedge
dashed wedge
When drawing 3-D diagrams of molecules two rules must be kept in mind:
• If there are multiple center atoms, try to put as many bonds (atoms) as possible in the plane of the
paper. Pick out a “row” of atoms that goes through the whole molecule and let it be in the plane
of the paper; draw it with the correct angles, then add the other atoms to be backward and
forward.
o Example: For butane
H
H
H
H
H
C
C
C
C
H
H
H
H
H
Select all the carbon atoms and one hydrogen atom from each end and draw it as follows:
H
C
C
C
C
H
Next, add the forward and backward atoms as shown in the following structure:
H
H
H
H
C
C
C
H
H
C
H
H
H
H
This is commonly drawn without showing any carbons:
H
H
H
H
H
H
H
H
.
• The bonds should be shown at approximately the correct angle in any size molecule. Keep in
mind that if there are several tetrahedral arrangements in a row, the forward and backwards
pointing bonds in a tetrahedral arrangement on the same atom are often shown to both point
upwards or both down, not one up and one down. If there are several tetrahedral arrangements
in a row, the up/down arrangement alternates. For example:
H
H
H
H
H
H
H
H
H
H
H
NOT
H
H
H
H
H
The same viewing angle should work for all molecules, and in the second drawing the angle
shown for the first two carbons is not the same as that for the last.
17
Name ________________________________
Prelab
1. What is the prefix used for an organic compound where the longest carbon chain contains …
a. 2 carbon atoms _______________
b. 4 carbon atoms _______________
c. 6 carbon atoms _______________
d. 8 carbon atoms _______________
2. What is the name of an alkyl substituent with …
a. 1 carbon atom _______________
b. 3 carbon atoms _______________
c. 5 carbon atoms _______________
d. 7 carbon atoms _______________
3. What is the name of each of the following halogens as a substituent …
a. fluorine _______________
b. chlorine _______________
c. bromine _______________
18
Name ________________________________
Organic Structures & Nomenclature Lab
Report Sheet
Procedure
Complete each of the following tasks and submit your report sheet at the end of the lab period. Use
the provided model kits when directed. Note: These kits contain different colored balls for the atoms
and holes to connect atoms with bonds. It is not that important if you use atoms of a particular color for
particular atoms as long as the ball used has enough holes for each electron pair. The rigid sticks are
used for single bonds and the flexible sticks must be used to make the double and triple bonds (one stick
per pair of electrons).
Part A. Draw Lewis structures of the molecules whose formulas are below and provide the name of
each. Remember Lewis structures show lone pairs even on outer atoms. Next, convert each Lewis
structure to a line structures. Note: In a line structure carbon and hydrogen atoms are not shown; each
end point of a line represents a carbon atom and each carbon is assumed to have enough hydrogen atoms
to fill carbon’s octet. Lastly, make models of each using the provided model kit. Using the model as a
guide, draw a 3-dimensional picture of each. Be sure to draw your structure showing the correct bond
angles!
Formula
CH3 Cl
C2 H5 Cl
Lewis
Structure
Name
Line
Structure
3-D
Structure
19
C3 H8
Part B. The following are isomers of C 7 H16 .
1
2
3
4
Name the molecules represented above.
1. ______________________________
2. ______________________________
3. ______________________________
4. ______________________________
The test of whether two organic compounds are the same or different is whether the two models
can be superimposed exactly on each other. This is most easily tested using models. When you try to
superimpose the models, you may do any twisting or turning of the single bonds you wish, but you
may not break any "bonds” in the models.
Which of the molecules represented below is/are the same as #2? __________
Which of the molecules represented below is/are the same as #3? __________
Which of the molecules represented below is/are the same as #4? __________
20
Part C. There are four isomers with the formula C4 H9 Br. Make models of each one and complete the
table below. Note that the –Br is a "bromo" substituent.
Line Structure
Name
#1
#2
#3
#4
How are isomers different than isotopes?
21
Part D. Your instructor will assign you one of the four groups of compounds from the table below
(#1, #2, #3, or #4). Write the name of each structure in your group right below the given line
structures. Hint: NO 2 is a “nitro” substituent.
#1
#2
#3
#4
(a)
F
F
(b)
Cl
Cl
F
Cl
Cl
(c)
F
Br
Br
(d)
Br
(e)
22
Cl
Experiment 2: Gas Chromatography, Intermolecular Forces, & Boiling Point
Background
Mixtures of many types are separated by the use of gas chromatography (GC), also called gas-liquid
chromatography (GLC). This type of separation is based upon differences in boiling points as well as
the intermolecular attractions (intermolecular forces) between the mixture components and the liquid
that coats the walls of the column. The boiling point of a compound is related to the intermolecular
attractions between its molecules. Often it is important to consider differences in the mass of molecules
instead of simply the types of intermolecular forces, because molecules with larger masses are typically
larger and thus experience stronger London forces.
In GC, a tiny volume of sample is drawn into a syringe and then injected into the end of a heated
tube (called a column) that has a liquid coating on the walls. This liquid coating is called a stationary
phase. A gas (called the mobile phase) is flowing through the tube. The inside of the column on most
modern instruments is hollow, and the sample and the mobile phase pass through the hollow part. The
column is many meters long and has a diameter in the low micrometer range. As the sample moves
through the column, the different components in the mixture repeatedly adsorb onto the stationary phase
and then come out into the mobile phase. After a number of times of adsorption/desorption, the sample
comes out the end of the column and goes into the detector. The lower the boiling point of a component,
the less it interacts with the stationary phase and the more quickly it moves down the column and out
into the detector.
side of column
from port
to detector
liquid coating
(a) early
(b) later
side of column
to detector
quid coating
(b) later
As the molecules pass through the detector, there is a signal, resulting in a peak on the
chromatogram. Even if all the molecules are identical in composition, some of them pass through the
detector a tiny bit sooner than the others, so the peak grows grad ually and then decreases. Ideally, the
peak should be symmetrical. If the mixture is separated well
then each component has its own peak. The time it takes for
the top of the peak to go through the detector is called the
amount
retention time and is used to tell which of various possible
of signal
compounds is causing that peak (from a legal viewpoint it is
not enough to identify the compound, but if you know what
time
might be there as in our samples, then you can choose between
the possible components). The percent of each component can
retention
time
be calculated by dividing the area of its peak by the sum of the
area of all the peaks present and then multiplying by 100.
23
Prelab
Name ______________________________
Part 1: Gas Chromatography
Watch the following videos and then answer the question below
• https://www.youtube.com/watch?v=3AQ55RPVE_A
• https://www.youtube.com/watch?v=4Xaa9WdXVTM
o Note that the Kahn Academy video discusses a difference in mass of the molecules with
the same boiling point, but this is really differences in intermolecular attractions. Be
aware that there are some differences in different types of chromatographs, such as types
of detector, and whether you have to make a solution of the sample or use it pure.
1. What information is obtained from the retention time of a peak in a chromatogram?
2. What information is calculated from the area of the peaks in a chromatogram?
3. If your sample contains three compounds and there is no solvent, how many peaks should you
expect in the chromatogram?
4. What does volatile mean?
5. If a sample is made using 0.145 g of methanol and 0.554 g of ethanol,
(a) What percent of the total peak area is expected for the methanol peak?
𝑎𝑟𝑒𝑎 𝑜𝑓 𝑚𝑒𝑡ℎ𝑎𝑛𝑜𝑙 𝑝𝑒𝑎𝑘
𝑎𝑟𝑒𝑎 𝑜𝑓 𝑚𝑒𝑡ℎ𝑎𝑛𝑜𝑙 𝑝𝑒𝑎𝑘 + 𝑎𝑟𝑒𝑎 𝑜𝑓 𝑒𝑡ℎ𝑎𝑛𝑜𝑙 𝑝𝑒𝑎𝑘
× 100
Show work.
(b) Look up the boiling points of the methanol and the ethanol and tell which comes out first (of
the two). Explain.
24
Part 2: Intermolecular Forces & Boiling Point
Review organic functional groups and intermolecular forces in your lecture text (sections 23.2
and 11.1) and then answer the following questions.
1. What is the suffix used when naming compounds with each of the following functionalities?
a) a hydrocarbon with only carbon-carbon single bonds __________
b) a hydrocarbon with a carbon-carbon double bond __________
c) a hydrocarbon with a carbon-carbon triple bond __________
d) an alcohol __________
e) an ester __________
2. Provide the Lewis structure for each of the following functional groups (only the functional group,
not a specific molecule):
a) an alcohol
b) an ester
3. For each of the following compounds, draw line structures of two molecules and use dotted lines to
indicate what atoms of the two molecules are attracted the most, i.e., use dotted lines to show the
strongest intermolecular attraction between the molecules.
a) ethanol
b) ethyl ethanoate (ethyl acetate)
4. Identify what type(s) of intermolecular forces are present between molecules of:
a) a hydrocarbon
□ London dispersion
□ dipole-dipole
□ hydrogen bonding
b) an alcohol
□ London dispersion
□ dipole-dipole
□ hydrogen bonding
c) an ester
□ London dispersion
□ dipole-dipole
□ hydrogen bonding
5. Considering compounds with similar molar masses but a different functionality, which is likely to
have the …
a) lowest boiling point
□ hydrocarbon
□ alcohol
□ ester
b) highest boiling point
□ hydrocarbon
□ alcohol
□ ester
25
Procedure
Part 1: Gas Chromatography
Your lab instructor will demonstrate how to do an injection properly using a standard mixture of
four known alcohols. You will need the chromatogram of the standard to help identify the components
in your unknown mixture (this will be discussed in lab). Each group will be assigned an unknown
mixture and have a turn using the GC. When it is your turn, be very careful with the syringe so you do
not bend the plunger! The best way to prevent this is to not draw the plunger very far out, as you won’t
need to do so. Be sure there are no air bubbles in the liquid before you inject it! While your group is
awaiting your turn on the GC (and/or after your turn on the GC) you will work on Part 2.
Part 2: Intermolecular Forces & Boiling Point
Use the resources provided in lab to complete the following table; for each of the given compounds
identify the type of compound then provide the condensed structural formula, the molar mass, and the
boiling point.
Name
Type of
Compound
butanol
(butyl alcohol)
□ hydrocarbon
□ alcohol □ ester
ethyl ethanoate
(ethyl acetate)
□ hydrocarbon
□ alcohol □ ester
□ hydrocarbon
□ alcohol □ ester
□ hydrocarbon
□ alcohol □ ester
□ hydrocarbon
□ alcohol □ ester
□ hydrocarbon
□ alcohol □ ester
□ hydrocarbon
□ alcohol □ ester
□ hydrocarbon
□ alcohol □ ester
□ hydrocarbon
□ alcohol □ ester
□ hydrocarbon
□ alcohol □ ester
□ hydrocarbon
□ alcohol □ ester
□ hydrocarbon
□ alcohol □ ester
□ hydrocarbon
□ alcohol □ ester
□ hydrocarbon
□ alcohol □ ester
□ hydrocarbon
□ alcohol □ ester
heptane
octane
pentane
pentanol
propyl ethanoate
(propyl acetate)
butyl ethanoate
(butyl acetate)
hexanol
hexane
methyl ethanoate
(methyl acetate)
heptanol
pentyl ethanoate
(pentyl acetate)
octanol
nonane
Condensed Structural
Formula
26
Molar Mass
(g/mol)
Boiling
Point (°C)
Report
Part 1: Gas Chromatography
• Provide your chromatographs of the standard mixture and your unknown mixture with the peaks
labeled identifying which alcohol correspond to each peak.
• Calculate the percentage of each alcohol in your unknown mixture (show your work directly on
your unknown chromatograph).
Part 2: Intermolecular Forces & Boiling Point
• Transfer the name, type of compound, molar mass, and boiling point data found in lab into an
Excel spreadsheet.
o Organize the data by functional group, i.e., group all of the hydrocarbons together, all of
the alcohols together, and all of the esters together.
• Graph the boiling point vs. molar mass as three sets of data on one graph (one for hydrocarbons,
one for alcohols and one for esters).
o All of this data should go on one graph, but as three sets of data, not one set.
o If you do not know how to put multiple data sets on one graph, see the graphing
instructions at the beginning of the lab manual.
• Put a best fit line through each of the data sets (so you should have three lines) and display the
equation for each of the three lines.
o The unit on the slopes are the (units of y/units of x); in this case it should be the change
in boiling point per one g/mol change in the molar mass.
o Hint: The three slopes should help you answer the first three questions below; the molar
mass increases by ~14 g/mol for each added CH 2 group.
• Transfer your table(s) of data and the graph of boiling point vs. molar mass to a Word document.
• Answer the questions below in the same Word document and submit a copy of the document
along with the items from Part 1.
1. On average, how much is the boiling point of a hydrocarbon changed by adding one CH 2 group
(increasing the molar mass by 14 g/mol)? Hint: You must use the slope from your hydrocarbon data.
2. On average, how much is the boiling point of an ester changed by adding one CH 2 group?
3. On average, how much is the boiling point of an alcohol changed by adding one CH 2 group?
4. Within the range of the molar masses of the substances used in this experiment, does the trend of
boiling points follow what the textbook would predict (substances whose molecules can hydrogen
bond boil at higher temperatures than substances with polar molecules that cannot hydrogen bond,
which are in turn higher than substances with nonpolar molecules)? Briefly explain using your data
(examine your graph to compare alcohols, esters, and hydrocarbons that have similar molar masses).
27
Experiment 3: Net Ionic Equations
Introduction
In this exercise you will be combining solutions of ionic compounds, acids, and bases and observing
the results. Throughout the process you will make observations of the initial solutions, thoroughly mix
the two solutions, observe any immediate changes that occur, and observe the final state of the mixture.
Then you will identify the ions that were initially present in the two solutions that were mixed and write
the net ionic equation that is consistent with what you observed (or indicate if no reaction occurred).
Although you will not be required to write out the molecular equation and the ionic equation for
these reactions, it can be helpful in determining the net ionic equation (especially if you are new to or
haven’t had much practice writing out net ionic equations previously). You should review section 4.2
in your lecture textbook if you need a review of how to write molecular and ionic equations. You should
also have the solubility rules from your lecture textbook (Tables 4.2 and 4.3) memorized before coming
to lab so that you can successfully distinguish between products that should be soluble and products
that should be insoluble.
Example Data Analysis
The example below summarizes the kind of data analysis we expect. Note that you need to write
the formulas of the substances, not the 1B, etc.
Example
Solutions mixed:
0.1 M CoSO 4 and 0.1 M NaOH
Appearance of initial solutions:
The CoSO 4 solution was pink and the NaOH solution was colorless.
Appearance after mixing:
Upon mixing we could see a blue solid and the liquid was pink.
Starting solutions contain:
𝐶𝑜 2+ , 𝑆𝑂42− , 𝑁𝑎 +, 𝑂𝐻 −
Net Ionic Equation:
𝐶𝑜 2+ (𝑎𝑞) + 2 𝑂𝐻 − (𝑎𝑞) → 𝐶𝑜(𝑂𝐻)2 (𝑠)
Example Notes:
To determine the net ionic equation, you must be able to predict the products of the
reaction and identify any insoluble products (or other weak or non-electrolytes formed,
such as water) because these will be part of the net ionic equation along with any ions that
are necessary in forming them. Alternatively you can first predict the products and write
the reaction in the molecular form, then convert it to the complete ionic equation, and
finally cancel the spectator ions to get the net ionic equation. For the above example the
following are the corresponding molecular and ionic equations that would help lead you to
the net ionic equation shown above.
𝐶𝑜𝑆𝑂4 (𝑎𝑞) + 2 𝑁𝑎𝑂𝐻(𝑎𝑞) → 𝐶𝑜 (𝑂𝐻)(𝑠) + 𝑁𝑎2 𝑆𝑂4 (𝑎𝑞)
𝐶𝑜 2+ (𝑎𝑞) + 𝑆𝑂42− (𝑎𝑞) + 2 𝑁𝑎+ (𝑎𝑞) + 2 𝑂𝐻 −(𝑎𝑞) → 𝐶𝑜 (𝑂𝐻)2 (𝑠) + 2 𝑁𝑎+ (𝑎𝑞) + 𝑆𝑂42− (𝑎𝑞)
28
Name ______________________________
Prelab
Memorize the solubility rules provided in Table 4.2 and 4.3 found in section 4.2 of your lecture
ebook. At least for now you should know the main rules (the information on the left of each table),
within about the next month it will become important that you have both the main rules and exceptions
memorized. You should also review how to predict products of reactions between ionic compounds,
predict products of reactions between ionic compounds as well as acids and bases, and how to write
reaction equations in the molecular, ionic, and net ionic forms. All of these things can be found in
sections 4.2 and 4.3 of your lecture ebook. Utilize this information to answer the following prelab
questions.
_____ 1. Based on the solubility rules, which one of these compounds is insoluble in water?
A) NaCl
B) MgBr2
C) FeCl2
D) AgBr
E) ZnCl2
_____ 2. Based on the solubility rules, which one of these compounds is insoluble in water?
A) Na2 SO4
B) BaSO 4
C) CuSO 4
D) MgSO 4
E) Rb2 SO 4
_____ 3. Based on the solubility rules, which one of these compounds is soluble in water?
A) Hg2 Cl2
B) Na2 S
C) Ag2 CO 3
D) Ag2 S
E) BaCO 3
_____ 4. Based on the solubility rules, which one of these compounds is soluble in water?
A) AgBr
B) AgCl
C) Ag2 CO 3
D) AgNO 3
E) Ag2 S
_____ 5. If aqueous solutions of Na2 CO 3 and BaCl2 are mixed, which insoluble precipitate is formed?
A) Ba2 CO3
B) BaCO 3
C) NaCl
D) NaCl2
E) BaO
_____ 6. If aqueous solutions of Mg(C 2 H3 O2 )2 and LiOH are mixed, which insoluble precipitate is
formed?
A) LiC2 H3 O2
B) Li(C2 H3 O 2 )2
C) MgOH
D) Mg(OH)2
E) CH 3 OH
_____ 7. If aqueous solutions of Pb(NO 3 )2 and NaCl are mixed, which ions, if any, are spectator ions?
Pb(NO 3 )2 (aq) + 2 NaCl(aq) → PbCl2 (s) + 2 NaNO 3 (aq)
A) Pb2+(aq), Cl– (aq)
D) Na+(aq), Cl– (aq)
B) Na+(aq), NO 3 – (aq)
E) None are spectator ions.
C) Pb2+(aq), NO 3 – (aq)
_____ 8. Identify the net ionic equation for the following reaction:
2 NaCl(aq) + Hg2 (NO 3 )2 (aq) → 2 NaNO 3 (aq) + Hg2 Cl2 (s)
A) Na+(aq) + NO 3 – (aq) → NaNO 3 (aq)
B) Hg2 2+(aq) + 2 Cl– (aq) → Hg2 Cl2 (s)
C) 2 Hg+(aq) + 2 Cl– (aq) → 2HgCl(s)
D) Hg2 (NO 3 )2 (aq) → Hg2 2+(aq) + 2 NO 3 – (aq)
E) Hg2 2+(aq) → Hg2 (s)
29
_____ 9. What is the net ionic equation for the reaction that occurs when aqueous solutions of
Pb(NO 3 )2 and NH 4 Cl are mixed?
A) Pb(NO 3 )2 (aq) + 2 NH 4 Cl(aq) → 2 NH 4 NO3 (aq) + PbCl2 (s)
B) Pb2+(aq) + 2 Cl– (aq) → PbCl2 (s)
C) Pb2+(aq) + 2 NO 3 – (aq) + 2 NH 4 +(aq) + 2 Cl– (aq) → 2 NH 4 +(aq) + 2 NO 3 – (aq) + PbCl2 (s)
D) NH 4 +(aq) + NO 3 – (aq) → NH 4 NO 3 (s)
E) NH 4 +(aq) + NO 3 – (aq) → 2 NO(g) + 2 H 2 O(l)
_____ 10. Complete the following reaction and identify the Brønsted acid.
NaOH(aq) + HCl(aq) →
A) NaH(aq) + HOCl(aq); NaOH is the acid.
B) NaH(aq) + HOCl(aq); HCl is the acid.
C) NaCl(aq) + H2 O(l); NaOH is the acid.
D) NaCl(aq) + H 2 O(l); HCl is the acid.
E) NaCl(aq) + H 2 O(l); NaCl is the acid.
_____ 11. If aqueous solutions of Ba(OH)2 and HNO 3 are mixed, what products are formed?
A) BaN 2 (s) + H 2 O(l)
C) Ba(s) + H 2 (g) + NO 2 (g)
E) Ba3 N2 (s) + H 2 O(l)
B) Ba(NO 3 )2 (aq) + H 2 O(l)
D) Ba2 O(s) + NO 2 (g) + H 2 O(l)
_____ 12. Which salt is produced by the neutralization of hydrobromic acid with magnesium
hydroxide?
A) MgBr
B) Mg2 Br3
C) MgBr2
D) Mg3 Br2
E) Mg2 Br
D) HBr
E) HCl
_____ 13. Which of the following is a weak acid?
A) H 2 SO 4
B) HNO 3
C) HF
_____ 14. Which of the following is a strong acid?
A) H 3 PO 4
B) HNO 3
C) HF
D) CH 3 COOH
E) H 2 O
_____ 15. Which is the net ionic equation for the reaction between aqueous solutions of lithium
hydroxide and hydrobromic acid?
LiOH(aq) + HBr(aq) → H 2 O(l) + LiBr(aq)
A) LiOH(aq) → Li+(aq) + OH – (aq)
B) HBr(aq) → H +(aq) + Br– (aq)
C) H +(aq) + OH – (aq) → H 2 O(l)
D) Li+(aq) + Br– (aq) → LiBr(aq)
E) Li+(aq) + OH – (aq) + H +(aq) + Br– (aq) → H 2 O(l) + LiBr(aq)
30
Procedure
PRECAUTIONS: ALMOST ALL THE SOLUTIONS USED IN THIS EXERCISE ARE
CORROSIVE.
ANY SPILLS MUST BE DILUTED WITH WATER AND WASHED
COMPLETELY AWAY IMMEDIATELY. Mixtures containing nickel, copper, or barium must be
disposed of in the bottle labeled "Heavy Metal Recovery"
At the start of lab your instructor will assign you one of the following sets:
(a)
(b)
(c)
(d)
(e)
(f)
(g)
(h)
(i)
(j)
1A, 1C, 2B, 3F, 4D, 5C, 6G, 8I, 11E, 11J, 12J, 13I
1B, 1C, 2A, 3F, 4G, 6F, 8D, 10J, 11I, 12I, 13E, 13J
1A, 1D, 2C, 3B, 4G, 8F, 9G, 10I, 11J, 12J, 13E, 13I
1B, 1D, 2C, 3A, 4G, 8F, 9D, 10J, 11E, 11I, 12I, 13J
1A, 1F, 2D, 3C, 4B, 5I, 9C, 10G, 11E, 11J, 12J, 13I
1B, 1F, 2G, 3C, 4A, 5D, 6I, 10F, 11J, 12J, 13E, 13I
1A, 1G, 2B, 3D, 4C, 5F, 6G, 9J, 11I, 12I, 13E, 13J
1B, 1G, 2A, 3D, 4C, 6F, 8D, 10I, 11E, 11J, 12J, 13I
1A, 1F, 2F, 3B, 4G, 6C, 8G, 9J, 11E, 11I, 12I, 13J
1B, 1G, 2F, 3A, 4D, 8C, 9D, 10I, 11J, 12J, 13E, 13I
Each code in your assigned set corresponds to two chemicals (one for the number and one for the
letter; see the table below) that you will combine, observe, and identify the reaction that occurs (if any).
For example, the code 8J means that you should combine a 0.1 M solution of FeSO 4 with a 0.1 M NaOH
solution and carefully observed the result.
1. 0.1 M MgCl2
2. 0.1 M CaCl2
3. 0.1 M SrCl2
4. 0.1 M BaCl2
5. 0.1 M ZnSO 4
6. 0.1 M CoCl2
8. 0.1 M FeSO 4
9. 0.1 M CuSO 4
10. 0.1M MnSO 4
11. 0.1 M HCl
12. 0.1 M CH 3 COOH
13. 0.1 M HNO3
A. 0.1 M Na2SO 4
B. 0.1 M (NH 4)2SO 4
C. 0.1 M Na3PO4
D. 0.1 M Na2CO3
E. NH 3
F. 0.1 M K 3PO4
G. 0.1 M K 2CO3
I. 0.1 M KOH
J. 0.1 M NaOH
For each of your assigned sets you should obtain about 1 mL of each solution in two separate test
tubes. Then write down observations of the initial appearance of each solution. Pour one solution into
the other and gently thump the side of the tube near the bottom to mix the solutions. Write down initial
observations of the mixed solutions. Set the tube aside for 10 minutes to ensure the reaction is complete
and to allow any precipitate to settle. Write down final observations of the mixed solutions. Lastly, write
down the net ionic equation for the reaction that occurred (this should be consistent with your
observations). It may be helpful to have some scratch paper to first write out the molecular and ionic
forms of the reaction before simplifying it to a net ionic equation.
31
Note: You will be assigned some combinations of solutions that will result in an acid-base reactions.
There will not be any change in appearance that occurs as a result of these reactions, but that does
not mean that there was no reaction. If the mixture does not produce a precipitate you should
determine whether or not an acid-base reaction could have happened which would produce the
nonelectrolyte water and therefore have a net ionic equation.
Please note each of the following tips:
1. The test tubes you use (13  100 mm) hold almost exactly 10 mL. Measure by estimating a
fraction of the test tube volume, not by using graduated cylinders, etc.
2. You should never have more than 3 mL total in a test tube so that you can efficiently mix the
contents of a test tube without stirring (“thump” the tube).
3. When mixing liquids in this exercise, use around equal volumes of the two liquids, keeping the
total volume below 3 mL. Observe what happens when you mix them. Allow them to stand for
at least 10 minutes before you decide about the final state.
Fill out the included report sheets as you perform the lab; you will be expected to
submit your completed report sheets before you leave the lab. See your instructor
if you have problems
32
Name:
Lab Day and Time:
Solutions Mixed:
Appearances of Initial Solutions:
Appearance After Mixing and
Centrifuging (if Needed)
Ions Present:
Net Ionic Equation:
Solutions Mixed:
Appearances of Initial Solutions:
Appearance After Mixing and
Centrifuging (if Needed)
Ions Present:
Net Ionic Equation:
Solutions Mixed:
Appearances of Initial Solutions:
Appearance After Mixing and
Centrifuging (if Needed)
Ions Present:
Net Ionic Equation:
Solutions Mixed:
Appearances of Initial Solutions:
Appearance After Mixing and
Centrifuging (if Needed)
Ions Present:
Net Ionic Equation:
33
Name:
Lab Day and Time:
Solutions Mixed:
Appearances of Initial Solutions:
Appearance After Mixing and
Centrifuging (if Needed)
Ions Present:
Net Ionic Equation:
Solutions Mixed:
Appearances of Initial Solutions:
Appearance After Mixing and
Centrifuging (if Needed)
Ions Present:
Net Ionic Equation:
Solutions Mixed:
Appearances of Initial Solutions:
Appearance After Mixing and
Centrifuging (if Needed)
Ions Present:
Net Ionic Equation:
Solutions Mixed:
Appearances of Initial Solutions:
Appearance After Mixing and
Centrifuging (if Needed)
Ions Present:
Net Ionic Equation:
34
Name:
Lab Day and Time:
Solutions Mixed:
Appearances of Initial Solutions:
Appearance After Mixing and
Centrifuging (if Needed)
Ions Present:
Net Ionic Equation:
Solutions Mixed:
Appearances of Initial Solutions:
Appearance After Mixing and
Centrifuging (if Needed)
Ions Present:
Net Ionic Equation:
Solutions Mixed:
Appearances of Initial Solutions:
Appearance After Mixing and
Centrifuging (if Needed)
Ions Present:
Net Ionic Equation:
Solutions Mixed:
Appearances of Initial Solutions:
Appearance After Mixing and
Centrifuging (if Needed)
Ions Present:
Net Ionic Equation:
35
Experiment 4: Rate Law for Bleaching of Crystal Violet by NaOH
***You must bring a flash drive to lab to save your data files.***
Objective
The purpose of this experiment is to determine the rate law for the reaction of the compound crystal
violet (methyl violet) with 𝑂𝐻 − ion. To accomplish this you will need to find the reaction order with
respect to crystal violet, the reaction order with respect hydroxide ion, and the value for the rate constant.
Introduction
When hydroxide ion is added to a crystal violet solution, the solution becomes colorless as the
crystal violet reacts. You can follow the reaction by placing a reaction mixture in a spectrophotometer
and observing the change in absorbance with time. When carrying out this type of experiment you first
need to find the wavelength where the maximum absorbance occurs (λ max ). It is also important that you
find a concentration with an appropriate absorbance for the rate measurements (not too high for the
instrument and not so low that the error in the measurements is excessive). This is done by finding the
absorbance at λmax of a solution whose concentration is known, and since absorbance is proportional to
concentration the necessary scaling factor for the concentration can be determined from the ratio of the
absorbances.
Once absorbance versus time data has been collected it can be plotted to help determine the order
of a reaction with respect to the absorbing substance. Although this analysis is often done with
concentrations, absorbance can be substituted for concentration since they are proportional. Therefore,
when a plot of the absorbance versus time is linear the reaction is zero order, when a plot of
ln(absorbance) versus time is linear the reaction is first order, and when a plot of 1/absorbance versus
time is linear the reaction is second order. The plots below show you each case for data from a zero,
first, and second order reaction; notice that plots with a mismatched order have curvature, e.g., when
first and second order data are plotted as absorbance versus time the result is a curve whereas the zero
order data is linear.
36
Prelab
Consider the following data collected for the decomposition of substance A:
Time (min)
0
60.0
96.4
157.5
•
•
•
•
•
[A] (M)
0.50
0.40
0.35
0.28
Prepare the three plots necessary to test whether the data is zero, first, or second order, i.e., [A]
vs. t, ln[A] vs. t, and 1/[A] vs. t.
Add a trendline to each plot and show the equation and R 2 value on each plot.
Provide a copy of your plots.
Identify whether the reaction is zero, first, or second order.
Report the rate constant for this reaction (be sure to include units with the value).
37
Procedure
The stock solution of crystal violet that is provided for you is an appropriate concentration for the
absorbance measurements, so you will not need to dilute the crystal violet solution. Make note of the
concentration of the crystal violet solution provided. Set up and calibrate your spectrophotometer:
□ Plug in the LabQuest 2 device (LQ2).
□ Attach the LQ2 to the cube spectrometer
□ Power on the LQ2. After an initial startup screen, a screen with a red stripe that shows
USB:Abs should appear (this is the home screen). If you don’t see this ask for help.
□ At the upper right hand corner there should be a gray box with Mode:Full Spectrum, which
is what you need for this part. If you don’t see this ask for help.
□ Choose Sensors at the top left of the screen.
□ Choose Calibrate -> USB: Spectrometer near the bottom of the gray menu. A new screen
should now appear, indicating the time necessary for warm up.
□ While the spectrometer is warming up, place a clean cuvette filled about 2/3 full with DI
water (a blank) into the cell holder of the spectrometer.
□ Once the spectrometer is warmed up and your blank is in place, choose Finish Calibration.
□ When the calibration is complete, choose OK to go to the original screen.
Obtain the full absorbance spectrum of the crystal violet solution from 400-800 nm and save the
spectrum:
□
□
□
□
Empty your cuvette and shake out the excess water.
Fill the cuvette about 2/3 full with the crystal violet solution.
Place the cuvette into the spectrometer.
On the LQ2 screen, tap the green triangle at the bottom left. Data acquisition should start
and the screen should change to show a full spectrum of your sample.
□ When the black line with circles stops changing (this should happen almost instantly), tap
the small box at the left bottom with the red square to stop data collection.
□ Save this file.
o Select File in the upper left corner and then Save.
o Touch the field at the top with the stylus and then change ‘untitled’ to your last name
and a number (e.g. Miller1, Miller2, etc.; be sure you know what each file name
corresponds to).
o Tap Done and then Save in the bottom right corner.
□ Empty the cuvette, rinse with DI water, and shake out the excess water so it is ready for your
first kinetics trial.
Use the full absorbance spectrum to determine the wavelength of maximum absorbance, λ max , in this
wavelength range. Set up the spectrophotometer to collect absorbance versus time data at the λ max :
□ In the upper left corner, choose the icon that looks like a meter; this will take you to the
screen with the red stripe.
□ In the upper right corner, choose the mode square.
□ In the drop down menu at the top, choose time based.
□ Tap the field next to Duration, change the number to 1800, and tap Done. Then make sure
the unit is set to s for seconds.
□ Choose OK.
□ It will ask about your data and go ahead and select Discard since you already saved it.
□ This will take you back to the screen with the red stripe.
□ Tap the red stripe and select Change Wavelength.
□ Input the λmax into the Selected Wavelength field and then tap OK.
38
Now, you are ready to perform your first kinetics trial. Accurately pipet a small volume of the diluted
crystal violet stock solution into a small, dry beaker (2-3 mL is sufficient). Pipet the same number of
milliliters of 0.5 M NaOH into another small, dry beaker. Rapidly pour one of the beaker’s contents into
the other, quickly but carefully swirl the contents to mix, and then quickly fill your cuvette about 2/3
full with the mixture using a transfer pipet, immediately place the cuvette in the spectrophot ometer and
tap the green triangle at the bottom left of the screen to start collecting data. If the initial absorbance is
less than 0.5 (you can read this from the live plot that is displayed or look at the absorbance reading
shown near the top right of the screen), then you did not do the mixing, transfer, and start of data
collection quick enough and you must start the trial over. Once you have successfully started a trial that
has an initial absorbance of 0.5 or greater you must allow the data collection to continue until the
absorbance drops below one tenth of its initial absorbance, e.g., if the initial absorbance is 0.5 then you
must continue data collection until the absorbance drops to 0.05 or if the initial absorbance is 0.7 then
you need to continue until the absorbance drops to 0.07, etc. Tap the button with the red square to stop
the collection early if the initial absorbance was too low or once the absorbance has dropped to below
one tenth of the initial absorbance. If you have to start a new trial it will ask you about discarding the
data after you tap the green triangle to start a new collection, in this case you should tap Discard because
you don’t want to add your new trial to the previous one. Once you have successfully collected data that
meets these criteria stop the collection and save the data.
Next you need to perform three more kinetics trials with diluted concentrations of NaOH (something
less than 0.5 M), saving each data set after it is collected. You get to choose the diluted concentrations
of NaOH that you will use, but your lowest should be somewhere in the range of 0.05-0.1 M and your
four NaOH concentrations (including the 0.5 M) should be equally spaced (or close to equally spaced).
Follow the same steps as the previous trial for each of these trials: tap the green triangle button to st art
data collection once you have inserted the cuvette, discard the previous data that is already saved, tap
the red square button to stop data collection, and save the data. Note: All of the trials must meet the same
absorbance criteria as the first (initial absorbance of at least 0.5 and final absorbance that is lower
than one tenth of the initial absorbance).
Lastly, you need to run a duplicate of one of your previous trials and save the data; if there is a trial
you feel you may have made an error on then you should duplicate that trial. Follow the same steps as
before. When you are done you will have five trials where four will have different NaOH concentrations
and one will be a repeat. Export your saved absorbance vs. wavelength data set and your five absorbance
vs. time data sets to a USB drive:
□
□
□
□
□
□
□
□
□
Unplug the spectrometer from the USB port on the LQ2.
Plug your USB drive into the USB port on the LQ2.
Choose File -> Open.
In the upper left hand corner choose the farthest left icon (this will show you the LQ2 file
directory instead of your USB file directory).
Choose the file you want to save (you have to open the file in order to be able to save it to
your USB). This will open the file, but it will not go to the screen that shows the graph
(while you could go back to the graph, it is not necessary to export the file).
Choose File -> Export (NOT Save).
Type a name for your file (it does not automatically pull the name that you gave it on LQ2).
Tap Done, then OK. Now the data has been transferred as a text file to your USB drive.
Repeat the File -> Open, File -> Export process for each of your data files.
Either in lab or as soon as you get to a computer, check that you can import your data into Excel.
DO NOT wait until the next week to check, otherwise if you cannot open your data then you may not
have time to get it so you can complete your lab report. Turn off the LQ2, unplug everything and put it
back how you found it.
39
Data Analysis
□ The end goal is to determine the rate law for the reaction, including the reaction order with
respect to crystal violet, the reaction order with respect to hydroxide ion, and the rate
constant for the reaction, i.e., to find m, n, and k in rate = k[crystal violet]m[OH⁻]n .
□ Find the reaction order with respect to crystal violet, m.
o Each kinetics trial needs to be analyzed to determine if the reaction is 0, 1 st, or 2nd
order with respect to crystal violet.
o For each kinetics trial, make a plot of absorbance vs. time, a plot of ln(absorbance)
vs. time, and a plot of 1/absorbance vs. time to test for 0, 1 st, or 2nd order, respectively.
▪ You should end up with three plots for each of your five kinetics trials, giving
you a total of fifteen plots.
▪ Note: Normally concentration is used to do this, but since concentration and
absorbance are directly proportional the absorbance can be used in place of
concentration; the same graph will be a straight line if we use absorbance
instead of concentration.
□ Find the apparent rate constant, k’, for each trial, where k’ = k[OH⁻]n .
o You can’t find the actual rate constant, k, directly because the reaction order of the
hydroxide ion is not yet known, therefore first you must find the apparent rate
constant, k’.
o Look at your plots for each run, and decide which is the straightest for the majority
of runs. This indicates whether the order with respect to crystal violet is 0, 1 st, or 2nd
order.
o When you have decided on the most linear set, use that set to determine k’ for each
trial. Hint: The rate constant is always equal to the magnitude of the slope, so k’ is
the absolute value of the slope of the linear plots.
o Note: Since you are using absorbance values instead of concentration, the rate
constant will be in terms of the absorbance instead of the concentration.
□ Find the reaction order with respect to [OH⁻], n, and the rate constant, k.
o Make a table of k’ and the corresponding [OH⁻] values.
o Use the values from the above table to make a second table with ln(k’) and ln[OH⁻].
o Make a plot of ln(k’) vs. ln[OH⁻], and add a trendline and its equation to the plot to
find n and k.
▪ Recall that k’ = k[OH⁻]n .
▪ By taking the ln of the above, it follows that ln(k’) = n ln[OH⁻] + ln(k).
▪ Thus, n is the slope of the plot and k can be calculated from the y-intercept
of the plot (k = ey-intercept).
40
Report
Your report must include each of the following items in this order. Your instructor will notify you
of any additional report expectations.
•
•
•
•
•
•
•
•
•
•
Solution data: concentration of crystal violet solution, concentrations of hydroxide ion for each
trial, and sample calculation for the dilutions.
The absorbance vs. wavelength plot and a statement of the λ max determined from it.
Kinetic plots for each of the five trials.
o These should be presented on five different pages, with the three plots for a given trial
on its own page so they can be compared easily.
o Each page must be clearly labeled with the trial’s hydroxide concentration.
o Each page should include a trendline and its equation on only the most linear plot for
that trial.
An explicit statement of the order of the crystal violet and how it was determined.
A table of k’ vs. [OH⁻] and a table of ln(k’) vs. ln[OH⁻].
The graph of ln(k’) vs. ln[OH⁻] with the trendline and its equation.
An explicit statement of the order of the hydroxide ion and how it was determined.
The calculation of numerical value of k.
A conclusion stating the rate law in as explicit a form as you can (meaning with the actual values
for k, m, and n).
A conclusion about the uncertainty in your measurements. Hint: The k’ for the runs that you
duplicated should help in estimating this uncertainty.
41
Experiment 5: Le Chatelier’s Principle
Purpose
This exercise is intended to give you first-hand experience observing chemical equilibria, how
changing intensive variables can affect the direction of the reaction, and how the latter relates to Le
Chatelier’s Principle (LCP).
Introduction
LCP states that if you start with a system at equilibrium and change one of its intensive properties,
then the equilibrium will shift to offset that change in intensive property. As a result of that shift, we
may see other changes. In this experiment, you will often use these other observed changes to deduce
which way the equilibrium shifted. In some chemical reactions, the product will have a different color
than the reactant. In this case, if the equilibrium is disturbed, a resulting color that is more product-like
indicates a higher concentration of product and thus that the equilibrium shifted to the right. If the color
becomes more reactant-like, then there is a higher concentration of reactant and the equilibrium must
have shifted left. Similar changes in equilibrium position are indicated by solids forming or dissolving.
Experiment Hints
•
•
•
•
•
•
•
Look down through the top of the test tube to observe color changes; observing this way will
minimize visual changes due only to dilution. To see that this is true, consider if you doubled
the volume of a solution by adding water and nothing else happened. Then looking from the top
of the tube would be looking through twice as much solution that is half as dark, which would
look the same color intensity as the original.
If an equilibrium shifts due to a dilution, it is NOT due to the change in concentration of
water! In this case, you will need to consider the total solute concentration and its changes
as the intensive property.
If a reaction is endothermic, heat may be regarded as a reactant and if a reaction is exothermic,
heat is a product. For this situation, an increase in temperature increases the “concentration” of
the heat.
Write careful notes about what you see for each part of each of the experiments on the provided
report sheets. At each step, including before mixing reagents:
o Report the colors observed. Your observations of changes during dropwise additions
should be summarized rather than reported “drop by drop”.
o Clearly indicate what your observations imply about species present and what the
changes in species present imply about the direction of the equilibrium shift of the main
reaction equation (ME).
o Trace changes from the added reagent to the changes in intensive properties at each step
of the way until you reach the LCP predicted shift in the main equilibrium.
o Do not skip steps!
Note that LCP cannot be applied until an equilibrium has been established and sometimes you
must mix things to establish the equilibrium before disturbing it.
Be sure it is absolutely clear which are observations and which are deductions based on the
observations and the equation.
Be careful with each and every section; proper use of terminology is extremely important! Do
not say things like “LCP predicts if the reactant concentration increases then the equilibrium
will shift to produce more product”. LCP does not say anything about making product—it
simply says the equilibrium would shift to decrease the concentration of the added reactant.
42
Example Report Format
Main chemical equilibrium equation:
Useful reaction equation:
𝐵𝑎2+ (𝑎𝑞) + 𝐶𝑟𝑂42− (𝑎𝑞) → 𝐵𝑎𝐶𝑟𝑂4 (𝑠)
yellow
yellow ppt.
2 𝐶𝑟𝑂42− (𝑎𝑞) + 2 𝐻+ (𝑎𝑞) → 𝐶𝑟2 𝑂72− (𝑎𝑞) + 𝐻2 𝑂(𝑙)
yellow
orange
A. Addition of Na2 CrO 4 to BaCl2 :
• Observations:
o Establishing Equilibrium: The yellow sodium chromate and the colorless barium
chloride solutions mixed to result in a yellow solution with a small amount of yellow
precipitate.
o As more sodium chromate solution was added to barium chloride solution, more
yellow precipitate was formed in a still yellow solution.
• Interpretation (direction of equilibrium shift):
o Since the product in the main equilibrium (ME) is a precipitate, the ME must have
shifted right.
• LCP Explanation:
o LCP would predict this because the concentration of chromate ion was increased due to
added reagent, so the equilibrium shifted right to decrease the chromate ion molarity.
B. Addition of 6 M HCl to the above mixture:
• Observations:
o As the colorless HCl was added to the above (A) mixture, the solid dissolved and the
solution turned from yellow to orange.
• Interpretation:
o Because the solid dissolved and something orange formed, the main equilibrium must
have shifted left.
• LCP Explanation:
o As HCl was added, the [H +] reacted with the CrO 4 2-, so its concentration decreased.
This decrease makes the main reaction equilibrium shift to the left to try to increase the
[CrO 4 2-], and a left shift results in less BaCrO 4 precipitate.
Note that in the above:
• In part A, there was no need to discuss the useful reaction equation as it was not useful in this
case. Only use the “useful equation” if you need it.
• Reactants are always described.
• The equilibrium shift of the main equilibrium (ME) is explained with LCP.
• LCP and equilibrium shifts are connected with observations
• LCP explanations are often worded as “The equilibrium shifted (fill in) in order to (increase,
decrease) the concentration of (fill in).
43
Prelab
Name: ______________________________
Review section 15.5 of your lecture textbook which introduces Le Chatelier’s principle. You should
be able to predict what will happen as a result of stress applied to a reaction at equilibrium.
_____ 1. For the following reaction at equilibrium, which choice gives a change that will shift the
position of equilibrium to favor formation of more products?
2 NOBr(g) ↔ 2 NO(g) + Br2 (g)
A)
B)
C)
D)
E)
ΔHºrxn = 30 kJ/mol
Increase the total pressure by decreasing the volume
Add more NO
Remove Br2
Lower the temperature
Remove NOBr selectively
_____ 2. For the following reaction at equilibrium in a reaction vessel, which one of these changes
would cause the I 2 concentration to increase?
2 NOI(g) ↔ 2 NO(g) + I 2 (g)
A)
B)
C)
D)
E)
ΔHºrxn = 45.3 kJ/mol
Add more NOI
Add more NO
Increase the pressure
Compress the gas mixture into a smaller volume
Decrease the temperature
_____ 3. The following reaction is at equilibrium in a sealed container.
N2 (g) + 3 H 2 (g) ↔ 2 NH 3 (g)
ΔH°rxn < 0
Which, if any, of the following actions will increase the value of the equilibrium constant, K c?
A)
B)
C)
D)
E)
Adding more NH 3
Adding more N 2
Increasing the pressure
Lowering the temperature
Adding a catalyst
_____ 4. The reaction system POBr3 (g) ↔ POBr(g) + Br2 (g) is at equilibrium. Which of the following
statements describes the behavior of the system if POBr is added to the container?
A) POBr will be consumed in order to establish a new equilibrium.
B) The partial pressures of POBr3 and POBr will remain steady while the partial pressure of
bromine increases.
C) The partial pressure of bromine will increase while the partial pressure of POBr decreases.
D) The partial pressure of bromine remains steady while the partial pressures of POBr 3 and
POBr increase.
E) The forward reaction will proceed to establish equilibrium.
44
_____ 5. At 450°C, tert-butyl alcohol decomposes into water and isobutene.
(CH 3 )3 COH(g) ↔ (CH 3 )2 CCH2 (g) + H 2 O(g)
A reaction vessel contains these compounds at equilibrium. What will happen if the volume of
the container is reduced by 50% at constant temperature?
A)
B)
C)
D)
E)
The forward reaction will proceed in order to reestablish equilibrium.
The reverse reaction will proceed in order to reestablish equilibrium.
No change occurs.
The equilibrium constant will increase.
The equilibrium constant will decrease.
_____ 6. When the substances in the equation below are at equilibrium, at pressure, P, and
temperature, T, the equilibrium can be shifted to favor the products by
CuO(s) + H 2 (g) ↔ H 2 O(g) + Cu(s) ΔHºrxn = –2.0 kJ/mol
A)
B)
C)
D)
E)
increasing the pressure by means of a moving piston at constant T.
increasing the pressure by adding an inert gas such as nitrogen.
decreasing the temperature.
allowing some gases to escape at constant P and T.
adding a more copper.
_____ 7. The reaction 2 H 2 O2 (g) ↔ 2 H 2 O(g) + O 2 (g) is exothermic, ΔHºrxn = –210 kJ/mol. Which
one of the following is correct?
A)
B)
C)
D)
E)
K P at 800 K is smaller than K P at 1200 K.
Temperature does not affect K P
K P depends only on the pressure.
K P at 1200 K is smaller than K P at 800 K.
K P depends on total pressure as well as temperature.
45
Name: ______________________________
Procedure
You will establish an equilibrium for three reactions, make changes to disturb the equilibria, observe
the impact of the disturbances, deduce what shift in the equilibrium occurred, and explain how LCP
explains the observations. The specific steps you need to follow and the questions you are expected to
answer for each reaction are outlined in the report form below. Fill out the included report form as
you perform the lab; it is due at the end of the lab session.
NaOH, HCl, HNO 3 and NH 3 ARE CORROSIVE. THEY MUST BE WASHED OFF
IMMEDIATLEY WITH LARGE QUANTITIES OF WATER. THEY ARE VERY
DAMAGING TO EYES! KEEP YOUR SAFETY EYEWARE ON AT ALL TIMES.
Part 1
Bromothymol blue is an acid base indicator. Acid-base indicators are large organic molecules that
can gain or lose hydrogen ions to form substances that have different colors. The protonated form is
abbreviated below as HIn and the deprotonated form is abbreviated as In⁻.
Main chemical equilibrium equation: HIn(aq) ↔ In⁻(aq) + H 3 O +(aq)
yellow
blue
Sometimes useful reaction equation: OH⁻(aq) + H 3 O +(aq) ↔ 2 H2 O(l)
1. Establish an equilibrium by adding 4-6 drops of bromothymol blue to ~3 mL of pH 7 buffer
solution in a small test tube.
a. Observations:
46
2. Now add 1 M HCl dropwise, until you see a distinct change, to the tube from step 1, to disturb the
equilibrium.
a. Observations:
b. Which direction did the ME shift? What intensive variable in ME changed due to what you did?
Did it increase or decrease?
c. LCP Explanation of how the ME shifted due to change in intensive variable in ME:
3. Add 1 M NaOH dropwise until you see a distinct change.
a. Observations:
b. Which direction did the ME shift? What intensive variable in ME changed that led to the ME
shift? Did it increase or decrease? How did the change that you made lead to the change in
intensive variable in ME?
c. LCP explanation of how the ME shifted due to change in intensive variable in ME
47
Part 2
Main chemical equilibrium equation: Co2+(aq) + 4 Cl−(aq) ↔ CoCl4 2-(aq)
pale pink
bright blue
1. Establish an equilibrium by adding 12 M HCl, dropwise, to 1 mL of 1 M CoCl2 until the solution
appears purple. Be patient!
a. Observations:
2. Now visualize cooling the sample and answer the following:
a. If the H rxn is positive I can treat heat as a ______________ and the equilibrium would shift
___________ to make more heat, therefore, if I cool the sample I would see:
b. If the H rxn is negative I can treat heat as a ______________ and the equilibrium would shift
___________ to make more heat, therefore, if I cool the sample I would see:
3. Now place the tube in an ice bath and allow it to stand until no further change occurs.
a. Observations:
Hrxn is ________________. I can tell because:
Carefully and thoroughly explain how you can use the color changes to decide whether Hrxn
is positive or negative. (Hint: what does the equilibrium shift tell you about whether heat is a
reactant or product?)
48
Part 3
Main chemical equilibrium equation: Mg2+(aq) + 2 OH −(aq) ↔ Mg(OH)2 (s)
Useful Reaction Equation 1: NH 3 (aq) + H 2 O(l) ↔ NH4 +(aq) + OH −(aq)
Useful Reaction Equation 2: H 3 O+(aq) + OH −(aq) ↔ 2 H 2 O(l)
Note: You will not need any more than one of the useful reactions at a time.
1. Establish an equilibrium by adding 1 drop of 1 M NH3 to 1 mL of 0.1 M MgSO 4 in a test tube.
Observe carefully. If a very light precipitate does not form then add another drop of NH 3 and
continue until a precipitate is seen (it may be only a little).
2. Add 1 M NH 4 Cl dropwise with shaking until no further change occurs.
a. Observations:
b. Which direction did the ME shift? What intensive variable in ME changed that led to the ME
shift? Did it increase or decrease? How did the change that you made lead to the change in
intensive variable in ME?
c. LCP Explanation of how the ME shifted due to change in intensive variable in ME:
49
3. Add 6 M NaOH dropwise with shaking until no further change occurs. (Do not add more than 6
drops)
a. Observations:
b. Which direction did the ME shift? What intensive variable in ME changed that led to the ME
shift? Did it increase or decrease?
c. LCP Explanation of how the ME shifted due to change in intensive variable in ME:
4. Add 6 M HNO 3 dropwise with shaking until no further change occurs.
a. Observations:
b. Which direction did the ME shift? What intensive variable in ME changed that led to the ME
shift? Did it increase or decrease? How did the change that you made lead to the change in
intensive variable in ME?
c. LCP Explanation of how the ME shifted due to change in intensive variable in ME:
50
Experiment 6: Chemiluminescence – Oxidation of Luminol1
Lab Practical
Introduction
Chemiluminescence is the production of visible light by a chemical reaction without production of
heat. It occurs when a reaction produces a molecule in an electronically excited state and that molecule
then releases a photon of visible light as it returns to the ground state. Commercially available lightsticks
are an example of chemiluminescence. A solution containing luminol is sometimes used in combination
with a solution of hydrogen peroxide by crime scene investigators to detect the presence of hidden blood
spatters. In this lab you will prepare a solution containing luminol and mix it with a hydrogen peroxide
solution and observe the resulting chemiluminescence.
Prelab
Completion of the prelab is your “ticket” to get into lab. You will not be allowed to work without
having these answered, and if you have to answer them at the beginning of the lab period, you will not
be allowed extra time on the lab itself. Your instructor may choose to give you a quiz over the proper
use of quantitative glassware and/or calculations similar to those needed to prepare solutions A and B
at the start of lab.
1. Show the calculation of how much ammonium carbonate monohydrate, (NH 4 )2 CO3 ∙H2 O, is needed
to make 100.0 mL of 2.5 M ammonium carbonate, (NH 4 )2 CO3 . Be sure to write your work clear
including units and the correct number of significant figures.
2. Show the calculation of how many milliliters of 3.0% H 2 O 2 is needed to make 500.0 mL of 1.0%
H2O2.
Note: These calculations are similar to those you will need to do in lab; in lab you will need to do the
calculations for how to make solutions A and B without notes or other assistance. You can practice in
advance and ask you instructor to check your work during office hours. It would be a good idea to
review how to make solutions from solids and how to dilute solutions, including the calculations
associated with these processes, before coming to lab (see section 4.5 in your lecture textbook).
51
Hazards
All chemicals (including solids) used in this experiment are harmful if taken internally. A solution
of hydrogen peroxide is a topical antiseptic and cleaning agent. Rinse it off if you get it on you.
Solution Preparation Hints
One mole of a hydrate provides one mole of the contained salt, but also brings along some water
molecules. Thus, when determining the amount of hydrate needed, the mass of the contained water
molecules must be included in the molar mass calculation.
As you may recall from the oscillating reaction in CHEM 111, when you are preparing a solution
with more than one solute you simply calculate how many grams of each solute you need to make the
desired volume (in this case 250.0 mL) of the given concentration for each solute as though it was the
only one present, then put all the solutes together before adding solvent.
You should dissolve the solutes in a beaker first (just be sure you don’t exceed 250 mL). This will
allow you to more easily stir the solution to get the solutes to dissolve and also not clog a funnel or the
neck of the flask with solid. After they dissolve, transfer the solution to the flask, rinse the beaker and
pour it into the flask and dilute to the line.
Procedure
At the start of lab you will be directed to either calculate the amount of each solute necessary to
prepare solutions A and B (see solution information below) or to take a quiz with similar calculations.
After submitting your calculations or quiz prepare solutions A and B. Once you have solutions A and
B notify your instructor you are ready to carry out the reaction. The apparatus you need to mix the
solutions will already be set up for you. This will be done with the room lights dimmed to better see the
chemiluminescence. Slowly pour solutions A and B simultaneously into the funnel and carefully
observe what happens. Afterwards, flush the tygon tubing and sink with lots of water.
Solution A: Make a 250.0 mL solution that is:
• 0.0377 M sodium carbonate
• 1.13×10-3 M luminol, C8 H7 O2 N 3
• 0.286 M sodium hydrogen carbonate
• 4.38 × 10-3 M ammonium carbonate
o Using ammonium carbonate monohydrate, (NH 4 )2 CO3 ●H2 O
• 1.60 × 10-3 M copper sulfate
o Using copper sulfate pentahydrate, CuSO 4 ●5 H 2 O
Solution B: Make 250.0 mL of a 0.15% H 2 O2 solution using the provided 1.875 % solution of H 2 O2 .
Report
Turn in anything requested by your instructor before you leave lab; this may include your
calculations if you did not turn them before making the solutions as well as observations.
1 B.Z.
Shakhashiri Chemical Demonstrations: A Handbook for Teachers of Chemistry Vol 1.
University of Wisconsin Press: Madison, WI.
52
Experiment 7: Spectrophotometric Determination of Copper in a Penny
Objective
The purpose of this experiment is to review how to prepare solutions and how to obtain and use
absorption spectra of solutions. You will also review how to use a volumetric flask and a volumetric
pipet to prepare a standard solution. You will use a spectrometer to obtain absorption spectra of your
standard solution and an unknown solution. This experiment will prepare you for the upcoming
Luminol Lab Practical where you will need to use volumetric flasks and volumetric pipets to makes
solutions.
Background
One of the most direct ways to measure concentration, especially for large numbers of samples, is
spectrophotometry. A basic spectrophotometer is shown in schematic diagram below. The light path is
shown by the long thick gray arrow.
The white light from a normal light bulb passes into the monochromator which discards all the light
except that for a single wavelength which passes through it. This sounds like a light filter and that is
just how a monochromator works. Indeed, glass optical filters can be used for this purpose if a fairly
wide range of wavelengths can be used. This light then passes through the sample cell and hits the
surface of a detector. The detector’s electrical signal then goes to the electronics (possibly including a
computer) where it is converted to an absorbance.
We can summarize all this by saying that monochromatic light is passed through the sample and a
reading related to the amount of light absorbed is given. The absorbance readings we use are calculated
as: absorbance = –log(I sample/I reference). The number is larger as the sample absorbs more light. The reason
we use absorbance to represent the amount of light absorbed by the sample is that absorbance should
be proportional to the concentration of the species absorbing light, that is:
𝑎𝑏𝑠𝑜𝑟𝑏𝑎𝑛𝑐𝑒 = 𝑘 × 𝑐𝑜𝑛𝑐𝑒𝑛𝑡𝑟𝑎𝑡𝑖𝑜𝑛
where k is a constant that depends on the nature of the species absorbing light, the wavelength used,
and the thickness of the sample. The size of the constant, k, is determined experimentally by making a
number of standards, which are solutions that contain different known concentrations of the species that
absorbs light. We then measure the absorbance of these solutions and make a graph of absorbance vs.
concentration which is called a calibration line. The unknown is then run and the calibration line can be
used to find the unknown concentration from its absorbance.
When doing spectrophotometry, we normally work at a wavelength where the species has maximum
absorbance (referred to as the 𝜆 𝑚𝑎𝑥 ), that is, where the species absorbs the most light. The 𝜆 𝑚𝑎𝑥 is
53
determined by making a solution of the colored species and measuring its absorbance as a function of
wavelength of visible light. There are spectrophotometers that do this; usually these are diode array
instruments and while they work in a way that is similar to the simple spectrophotometer shown above,
they produce a whole spectrum (absorbance at all wavelengths) instead of a single reading of absorbance
at a single wavelength. Shown below are the spectra of 3 different solutions of the same dye,
superimposed on each other.
Two points should be made about these spectra. The main point is that no matter what the concentration
is, the wavelength of the absorption maximum and general shape of the spectra are the same, just with
a different absorbance. This is what is expected if the only change is the amount of the colored material.
No matter how much variation of concentration there is, the shape of the spectrum remains the same
with just changes in height. Second, we can look at the spectrum and see that the absorbance maximum
is at 631 nm and the absorbance is around 1.6 for a 10 mg/L standard. Also note that the absorbance is
around 1/2 this value in the 5 mg/L standard and 1/4 this value for the 2.5 mg/L standard. This is enough
information to let us know that absorbance measurements for solutions of this substance should be made
at a wavelength of 631 nm. We should also be aware that a maximum absorbance of 0.8 is to be used
for optimal accurate measurements, so standards should be made that cover the range 0 – 5 mg/L.
Once standards are made and absorbances are measured, the concentration and absorbance data are
tabulated and plotted. Below is a table of example data along with the resulting calibration line.
The graph itself could then be used to determine the concentration of an unknown solution from its
absorbance, but it is more accurate to use the calibration line equation to determine the unknown's
concentration by plugging in its absorbance in for y and solving for x which is the concentration.
54
Prelab
Name: ______________________________
You must bring a penny in 1983 or later to lab.
Go to youtube.com to watch the videos on reading the meniscus, the use of a pipet, volumetric flask.
After watching Youtube, complete the following questions. These will be due at the beginning of lab.
• Reading the Meniscus: https://www.youtube.com/watch?v=QCzaPy_XOqg
• Use of a volumetric flask to make a solution:
https://www.youtube.com/watch?v=Btp1N5Z2L74
o Notes:
▪ He has already rinsed the pipet well. Invert it more times than he did.
• Use of a pipet: https://www.youtube.com/watch?v=TffTiRw8cQY
• Review how to do dilution calculations in your textbook (pp.158-9)
Using a Volumetric Flask to Make a Solution
1. George needs to make 250.0 mL of an aqueous solution. He weighs the appropriate amount of solute
out and then adds 250.0 mL of water to it, stirs to dissolve the solid and mixes it well. What did he
do wrong?
2. What should you do with the volumetric flask before placing solute or solute solution into it?
3. Which of the following shows the proper liquid level in a full volumetric flask that is being used for
quantitative purposes?
4. a) At what points should you mix the solute and solvent by shaking the flask well?
b) Why should you do this both before the liquid is filled to the line and then again after?
c) Why should the flask be inverted during some of this mixing?
5. What should you do with the flask when you finish with it?
55
Using a Pipet
1. Briefly describe the procedure used to rinse the pipet properly.
2. Where should the tip of the pipet be compared to the liquid surface …
(a) while liquid is being drawn into the pipet?
(b) while liquid is being let down to the line on the pipet stem?
3. Which of the following shows the proper adjustment of liquid level in a pipet so it will deliver its
stated volume?
4. What do you do about any droplets or partial droplets of liquid on the outside of the pipet?
5. Should any liquid should be left in the tip of the pipet after you have used it to transfer the labeled
volume of liquid or should you blow it out?
6. Where should you look to find the precision of the pipet (so you know how many decimal places to
record in the volume)?
7. What should be done with the pipet before it is returned to the pipet storage container?
You will need a USB drive to save you data when you are in lab, so be sure to bring
one with you to lab.
56
Procedure
Each student (or pair) will dissolve a penny in dilute nitric acid, and then boil this to
remove some of the excess nitric acid. The solution will then be made basic and diluted to
some definite volume in a volumetric flask. Each student (or pair) will then measure the
absorbance of this penny sample. The analyte is Cu2+ and the brightly colored species is
the royal blue complex ion of copper with ammonia formed by the reaction:
Cu2+ + 4 NH 3 → Cu(NH 3) 4 2+
pale aqua colorless
royal blue
.
Each student (or pair) in the class will also need to make a standard Cu2+ solution by diluting a lab
stock solution (found in a large plastic bottle on the center bench) after adding the standard amount of
ammonia to it. The absorbance of the standards will be measured and recorded (for the class data) for
the absorbance vs. concentration graph. Then students will be able to determine the amount of copper
in their penny from the class absorbance vs. concentration graph that they will construct after measuring
the absorbance of their digested penny solution.
NITRIC ACID (HNO3 ) AND 15 M AMMONIA ARE CORROSIVE. AVOID SKIN CONTACT;
WASH AWAY ANY ACID OR AMMONIA WITH LARGE AMOUNTS OF WATER. 15 M
AMMONIA (NH3 ) HAS A VERY STRONG ODOR AND MUST BE KEPT IN THE HOOD AT
ALL TIMES. IT MAY NOT BE REMOVED FOR ANY REASON. DO NOT CARRY OPEN
CONTAINERS OF 15 M AMMONIA IN THE LAB. USE EVERY REASONABLE
PRECAUTION TO PREVENT RELEASING AMMONIA VAPOR IN THE LAB. WASH ANY
AFFECTED AREA WITH LARGE QUANTITIES OF WATER.
Procedure for preparing penny solution:
Weigh and record the mass of a post-1983 penny. Put the penny into a 100-150 mL beaker and add
30-40 mL of 6 M HNO 3. Cover the top with a watch glass, and gently warm the mixture to dissolve
(digest) the penny. [CAUTION: IN HOOD]. After the penny is dissolved, boil away some of the excess
acid, cool the sample and add the entire solution to a 250 mL volumetric flask. Measure out 50 mL of
15 M NH 3 (the ammonia does not have to be pipeted) and use a medicine dropper and small portions
of it to rinse the beaker three times, putting the rinse into the flask. Put the rest of the ammonia into the
flask, and rinse the penny beaker with a small portion of DI water and put it into the flask (do not go
over the line). Fill the flask to the line with DI water and mix it well.
Procedure for preparing standards:
Each student will make 100.0 mL of a standard of assigned concentration in the range of about 20–800
mg Cu2+/L by mixing the (pipeted) volume of stock appropriate for their assigned concentration with
25 mL of 15 M ammonia in a 100 mL volumetric flask. The mixture should be shaken vigorously. They
should then record the spectrum of the 500 mg Cu 2+/L in NH 3 solution available on a diode array
spectrophotometer, noting the wavelength at which the absorbance is highest. This wavelength of
highest absorbance depends only on the identity of the copper + ammonia complex present and not on
concentration or amount of other colorless materials that are present.
57
Collecting absorption spectra. You will collect absorption spectra for each of your five standards
(test tubes 1-4 and the stock solution) as well as an unknown provided by your instructor. Follow the
instructions below to collect and save your absorption spectra.
1. First set up your spectrometer using the following steps.
a. Plug in the LabQuest 2 device (LQ2).
b. Attach the LQ2 to the cube spectrometer
c. Power on the LQ2. After an initial startup screen, a screen with a red stripe that shows
USB:Abs should appear (this is the home screen). If you don’t see this ask for help.
d. At the upper right hand corner there should be a gray box with Mode:Full Spectrum,
which is what you need for this experiment. If you don’t see this ask for help.
e. Use the stylus attached to the cord (you likely need to pull it out of its storage slot on
the back of the spectrometer) to choose Sensors at the top left of the screen.
f. Choose Calibrate -> USB: Spectrometer near the bottom of the gray menu. A new
screen should now appear, indicating the time necessary for warm up.
g. While the spectrometer is warming up, place a clean cuvette filled about 2/3 full with
DI water (a blank) into the cell holder of the spectrometer. Note:
i. The spectrometer must be placed on a flat surface to hold the cuvette in place.
ii. Some of the cuvettes are stained or have other imperfections from prior use,
so it is best to use the same cuvette for the entire experiment because this
blanking process will account for these things in your spectra.
iii. When you put the liquid into a cuvette be sure there are not air bubbles; this
can happen when you put the liquid in too quickly.
iv. Only touch the cuvette on the frosted sides and wipe the clear sides with a
Kimwipe to remove moisture and fingerprints just before inserting it into the
spectrophotometer. You do not need a new Kimwipe everytime you do this,
and be sure to keep loose Kimwipes out of the hood!
v. The clear sides of the cuvette should be oriented so the light beam is travelling
through them.
h. Once the spectrometer is warmed up and your blank is in place, choose Finish
Calibration.
i. When the calibration is complete, choose OK to go to the original screen.
2. Collect an absorbance vs. wavelength spectrum for each of your five standards and your
unknown using the following steps.
a. Empty your cuvette and rinse it with a small amount of the solution you will be
measuring.
b. Fill the cuvette about 2/3 full with the solution to be measured.
c. Place the cuvette into the spectrometer.
d. On the LQ2 screen, tap the green triangle at the bottom left. Data acquisition should
start and the screen should change to show a full spectrum of your sample.
e. When the black line with circles stops changing (this should happen almost
instantly), tap the small box at the left bottom with the red square to stop data
collection.
f. Save this file.
i. Select File in the upper left corner and then Save.
ii. Touch the field at the top with the stylus and then change ‘untitled’ to your
last name and a number (e.g. Miller1, Miller2, etc.; be sure you know what
each file name corresponds to).
iii. Tap Done and then Save in the bottom right corner.
58
g. Repeat this process for each standard and your unknown. Note: When you start a
new data collection it will first ask if you want to discard the old data even if you
already saved the data; this will not eliminate your saved data it will just start a new
data collection.
3. Follow the steps below to export your data in text format so that you can process it in Excel.
a. Unplug the spectrometer from the USB port on the LQ2.
b. Plug your USB drive into the USB port on the LQ2.
c. Choose File -> Open.
d. In the upper left hand corner choose the farthest left icon (this will show you the LQ2
file directory instead of your USB file directory).
e. Choose the file you want to save (you have to open the file in order to be able to save
it to your USB). This will open the file, but it will not go to the screen that shows the
graph (while you could go back to the graph, it is not necessary to export the file).
f. Choose File -> Export (NOT Save).
g. Type a name for your file (it does not automatically pull the name that you gave it
on LQ2).
h. Tap Done, then OK. Now the data has been transferred as a text file to your USB
drive.
i. Repeat the File -> Open, File -> Export process for each of your data files.
4. Either in lab or as soon as you get to a computer, check that you can import your data into
Excel. DO NOT wait until the next week to check, otherwise if you cannot open your data
then you may not have time to get it so you can complete your lab report.
5. Turn off the LQ2, unplug everything and put it back how you found it.
59
Report
You should report your observations of colors and changes which occur during this experiment.
Tabulate the class standards data, make a calibration line from this data, and get an equation and R 2
value for the best straight line through the standards data. You may discard any data point you like in
making the curve but you must state which points were discarded and why in your final discussion.
Use the equation and show how to find the mg Cu in the penny and the % Cu in the penny.
Finally conclude how reliable this result is, using your results, and average of % Cu.
You should also compare your result with the stated value from the U.S. Department of the
Treasury, U.S. mint (usmint.gov) and calculate a % error. Don't hesitate to include a critique of the class
data for the calibration line if it looks awful or suspicious to you.
Also state which specific data points need to be omitted (if any), and why you would consider omitting
them.
*Your instructor will let you know of any further expectations.
60
Experiment 8: Synthesis of an Oxalate Hydrate
Introduction
One of the lab activities of organic, inorganic and pharmaceutical chemists is the synthesis of new
compounds. When they start a project, they have a target compound in mind, and after they isolate a
product they must find out if they have the desired product. This analysis may be done several ways,
but determining the mass percent composition is one step. In this experiment, you will perform the
synthesis step of this process, specifically, you will make a hydrate containing potassium, iron(III) and
oxalate ions. Your synthesized compound will have the generic formula K x Fey (C2 O4 )z∙QH 2O, where x,
y, z, and Q are unknown values. In the next experiment you will do a titration to determine the percent
of oxalate in your synthesized compound.
Procedure
•
First be aware of the following:
o You must rinse all sink surfaces well when pouring solutions containing iron down the
drain, as upon standing in air the iron will react to form rust stains.
o The final product will decompose upon long exposures to light, so store it in your desk
drawer when you are not using it.
•
Synthesis:
o In a flask, mix 30.0 mL of 1.95 M potassium oxalate and 10.0 mL of 1.5 M iron(III)
chloride and filter off the resulting solid.
Recrystallization (purification):
o Add about 5 mL hot water to an Erlenmeyer flask containing the crystals and heat as you
stir vigorously.
o If the crystals did not all dissolve, add about 0.5-1 mL of hot water and continue to stir
vigorously while heating.
▪ Repeat this process until the crystals dissolve.
▪ Caution: If there is a tiny amount of crystals that will not all dissolve do not keep
adding water as it is likely to be an impurity and adding too much water will cut
your yield! If this happens try to get the impurity out before the next step.
o After all the crystals dissolve remove the flask from the heat, cool the solution.
▪ Run water on the outside of the flask, but if it is too hot to hold with your hands
you will need to first let it cool in the air before running water over it.
▪ Next put the flask in an ice-water bath or add 10 mL of ethanol to ensure the
solution is cooled enough for maximum crystallization.
o Filter off the crystals formed.
o Spread the crystals on a piece of filter paper and leave them to dry in your lab drawer.
•
61
Report
You do not need to prepare a lab report for your synthesis reaction, instead you need to complet the
following literature search activity:
Literature Use: Finding a Cited Journal in the American Chemical Society (ACS) Web Editions
The journal format of ACS publications (other than Biochemistry) follow one of the following formats,
depending on which journal is used:
• Author 1; Author 2; etc. Title of Article. Journal Abbreviation Year, Volume, Inclusive
Pagination
• Author 1; Author 2; etc. Journal Abbreviation Year, Volume, Inclusive Pagination.
For example one citation is:
• Caruso, R. A.; Susha, A.; Caruso, F. Multilayered Titania, Silica, and Laponite Nanoparticle Coatings on
Polystyrene Colloidal Templates and Resulting Inorganic Hollow Spheres. Chem. Mater. 2001, 13, 400–
409.
o For this article, the following information would be needed to look it up using the ACS Web
Editions:
▪ Journal Name Abbreviation: Chem. Mater.
▪ Volume: 13
▪ Page: 400.
• If each issue starts with page one, the issue number may be in parentheses after the volume
number.
You will be assigned one or two of the citations below to look up. Go to MSU library home web page
http://www.moreheadstate.edu/library/. Under Other Search Options, choose ACS Web Editions, then
enter your MSU ID number and password. At right, toward the top, choose the Citation Tab and enter
the volume number and page number for your assigned article. Go to the PDF version of the full article
and print page one to turn in next week. Don’t print only the abstract—do full page 1.
1. Organometallics, 2004, 23 (15), pp 3562–
3583
3. Organometallics, 2011, 30 (20), pp 5338–
5343
5. J. Org. Chem., 2012, 77 (18), pp 7804–7814
7. J. Org. Chem., 2007, 72 (2), pp 313–322
9. Anal. Chem., 2013, 85 (24), pp 11677–
11680
11. J. Am. Chem. Soc., 2012, 134 (44), pp
18197–18200
13. J. Agric. Food Chem., 2014, 62 (51), pp
12418–12427
15. J. Phys. Chem. B, 2014, 118 (51), pp 14913–
14921
17. J. Am. Chem. Soc., 2006, 128 (48), pp
15354–15355
19. J. Am. Chem. Soc., 2014, 136 (51), pp
17692–17701
21. J. Med. Chem., 2016, 59 (23), pp 10435–
10450
62
2. J. Nat. Prod., 2013, 76 (9), pp 1541–1547
4. Anal. Chem., 2014, 86 (24), pp 11981–
11985
6. J. Am. Chem. Soc., 2012, 134 (51), pp
20581–20584
8. J. Nat. Prod., 2014, 77 (6), pp 1275–1279
10. J. Agric. Food Chem., 2017, 65 (2), pp 364–
372
12. J. Med. Chem., 2016, 59 (18), pp 8149–8167
14. J. Am. Chem. Soc., 2011, 133 (2), pp 200–
202
16. J. Med. Chem., 2007, 50 (22), pp 5269–5280
18. J. Nat. Prod., 2013, 76 (9), pp 1541–1547
20. J. Med. Chem., 2016, 59 (23), pp 10343–
10382
22. J. Am. Chem. Soc., 2013, 135 (51), pp
19079–19082
Experiment 9: Percent Oxalate Determination by Titration
You will do a titration of your oxalate hydrate from Experiment 9 using potassium permanganate in
order to determine the percentage of oxalate in your oxalate hydrate sample. The oxalate ion in your
hydrate reacts according to the following equation:
5 𝐶2 𝑂42− (𝑎𝑞) + 2 𝑀𝑛𝑂4− (𝑎𝑞) + 16 𝐻 + (𝑎𝑞) → 10 𝐶𝑂2 (𝑔) + 2 𝑀𝑛2+ (𝑎𝑞) + 8 𝐻2 𝑂(𝑙)
The slightly pink endpoint is reached when the first drop of excess permanganate is added.
Procedure
Grind about 1 gram of your product in a mortar and pestle until it forms a powder. Accurately weigh
0.15-0.20 grams of it into a titration flask and add 30 mL of DI water. Add 10 mL of 3 M sulfuric acid
and swirl the flask contents until the solid dissolves; if a brownish color develops, add more acid.
Use a buret to add about 10 mL of potassium permanganate to the titration flask (be sure you write
down your initial buret volume because you will need it to determine the total amount of titrant added
once the titration is complete). Add about 1-2 mL of 0.2 M MnCl2 and wait until the color of the KMnO4
disappears (the initial reaction is slow, but the manganese(II) formed catalyzes the reaction). Note: Mn2+
is a product of the reaction, so it is OK to add a little of it, it will not mess up the stoichiometry of the
titration. After the color disappears, continue the titration until the faint pink endpoint color persists for
20-30 seconds. At this point all of the oxalate ion has reacted, so record the final buret reading. Repeat
this titration two more times for a total of three trials (or repeat additional times for any that you mess
up). Use the total volume of titrant per trial to calculate the mass of oxalate in your titrated hydrate
sample and then calculate the mass percent of oxalate in your hydrate.
63
Experiment 10: Bleach Titration
Lab Practical
Objective
The purpose of this experiment is to determine the concentration of sodium hypochlorite in an
unknown bleach solution.
Introduction
Two of the most common types of reactions are acid -base reactions and oxidation-reduction (redox)
reactions. It is not surprising that such reactions are the basis for many analytical procedures used to
determine the amounts of substances in samples. An oxidation/reduction titration involves the addition
of an oxidizing agent to a solution of a reducing agent or vice versa. One of these solutions has a known
concentration while the other is unknown. The solutions are added until the stoichiometric amounts of
oxidizing and reducing agents have been combined.
The oxidizing agent in liquid bleach is sodium hypochlorite. To determine its concentration, it will
be added to a solution of potassium iodide to produce iodine, which is then titrated with a solution of
sodium thiosulfate. Thus, there are two reactions which will take place. First:
(1)
2 HCl(aq) + 2 KI(aq) + NaOCl(aq) → I 2 (aq) + NaCl(aq) + H 2 O(l) + 2 KCl(aq)
After the iodine is made in reaction (1), it is reduced back to the iodide ion by the sodium thiosulfate
as shown in the following equation (2).
(2)
I 2 (aq) + 2 Na2 S2 O3 (aq) → 2 NaI(aq) + Na2 S4 O6 (aq)
Prelab
You MUST do the prelab to start lab. Start early and if you do not know how, see your
instructor for help. Your instructor may choose to not let you in lab at all if you do not have the
prelab done correctly at the start of lab.
Have the following worked out carefully with proper unit cancellation shown. These will be checked
before you enter the lab! DO NOT use C1 V1 = C2 V2 (or M1 V1 = M2 V2 ) as that is for dilutions, not
titrations. Although this equation works in a limited number of titrations, it will not work for many of
them. This problem’s numbers are not the numbers you put in your report — it is simply an example so
you can be sure you know how to do the calculations in lab. You will have to do the calculations from
scratch in the lab without the prelab as a crutch, and you will need to turn in your calculations from the
titrations you do in lab before you leave. Part a is a scale-up calculation, and part b is a standard titration
calculation.
a. Suppose you did an exploratory titration and 2.00 mL of the bleach took 4.60 mL of Na 2 S2 O3 (aq).
How many mL of bleach would you use to end up needing 25-40 mL of Na2 S2 O 3 (aq)? Round
this to the nearest mL due to pipet sizes available. Show work. Note: There is a range of correct
answers, yours just needs to fall in the correct range based on the given criteria.
b. What is the molarity of NaOCl in a solution of bleach if 10.00 mL requires 22.35 mL of 0.02191
M Na2 S2 O 3 to reach the endpoint? Show work. Hint: Determine the amount of I 2 needed in the
second reaction, then use that to determine the amount of NaOCl in the first reaction. If you
don’t remember how to do titration calculations you may want to review section 4.6 from your
lecture textbook. Note: The correct answer is 0.02448 M.
64
Safety
Liquid bleach is corrosive and must be handled with care. Any spill on your body or in your eyes
must be washed with water immediately.
Suggestions
•
•
•
•
You should approach this lab like the acid-base titration lab except for the fact that you should
do these titrations slowly after you add the starch.
Don’t get in a hurry to get done. The color change after the addition of the starch is quick and if
you add titrant too quickly you will miss it.
Be sure to add the starch when the solution is a bright yellow like the outside of a lemon, do not
wait until it is light yellow or you will likely already be past the endpoint.
Note that it is the total amount of titrant per trial that is used to do the calculation — so take an
initial reading before beginning the titration, add titrant, add starch, finish the titration and take
a final reading. The difference in the two readings is the total volume of titrant used in the
titration and to be used in your calculation.
Lab Practical Overview
You will not be allowed to use any of your own copies of the lab or any additional notes. Your
instructor will provide you new copies of the lab instructions and blank paper for your results and
calculations. Each of you will be given a vial of unknown bleach solution. This unknown bleach solution
must be quantitatively transferred to a 100 mL volumetric flask and diluted to the mark (meaning you
need to transfer the entirety of the unknown solution into the volumetric flask; it is important that you
don’t leave any behind or lose any of it). The diluted solution in the flask is the solution whose NaOCl
concentration you will report, not the original unknown solution. You should begin by carrying out a
rapid course titration to determine the amount of diluted unknown bleach solution that will require about
20-40 mL of titrant (thiosulfate) after doing a scale-up. You only do the scale-up once! However, be
sure that you are not using more than 30 mL of your diluted bleach solution for each titration, otherwise
you will not have enough to do the minimum requirement of 3 fine titrations. You need to create a data
table of your results for each titration (including the course titration), show work for your scale-up
calculation, and show work for the calculation of the concentration of NaOCl in your bleach unknown
from your first fine titration on the provided data sheet that must be turned in at the end of lab.
Before each titration you should:
A. Have your buret filled with the titrant (sodium thiosulfate, Na2 S2 O3 ) and ready to go with the
initial volume reading written down.
B. Transfer your needed amount of diluted bleach solution into your Erlenmeyer flask using a pipet.
Note: You should only use a small volume for your course titration, then calculate how much
you will use for the for the fine titrations (the volume of diluted bleach solution should be large
enough that it requires about 20-40 mL of titrant, but not more than 30 mL of your bleach
solution otherwise you will not have enough for the minimum requirement of 3 fine titrations).
C. Have around 20 mL of DI water measured (it doesn’t need to be exact) and ready to add to the
Erlenmeyer flask with your measured diluted bleach solution.
D. Have about 1-1.5 g KI measured (also does not need to be exact) and ready to add to the
Erlenmeyer flask.
E. Have 10 mL of 1 M H 2 SO4 measured and ready to add to the Erlenmeyer flask.
65
Procedure
1. Quantitatively transfer your unknown bleach solution into a 100-mL volumetric flask and dilute
to the mark with DI water.
2. Pipet your needed volume of the diluted bleach solution into a 250-mL Erlenmeyer flask.
3. Add about 20 mL of DI water to the Erlenmeyer flask and stir.
4. Add about 1-1.5 g of potassium iodide to the Erlenmeyer flask and stir until it is dissolved.
5. Add 10 mL of 1 M H 2 SO4 ; this addition can be done with a graduated cylinder. The resulting
solution should be an orange, dark red, or brown color (this depends on your unknown bleach
concentration).
6. Titrate the sodium thiosulfate into the Erlenmeyer flask until the solution turns bright yellow.
Be sure you take an initial volume reading of the titrant before beginning the titration!
7. Add a full Pasteur pipet of starch solution to the Erlenmeyer flask. Upon the addition of starch,
your solution should turn a gray, blue, or almost black color (this depends on how close you are
to the endpoint and your total volume). If the solution is colorless after add the starch, then you
overshot the endpoint and will have to start the titration again.
8. Add titrant slowly until the solution just turns colorless. Take the final volume reading of the
titrant.
*You will have adequate precision when the volumes of titrant used for three trials in a row are
within 0.2 mL.
*If you violate safety procedures or do not clean up well, your instructor may count off.
Report
Prepare a table with all of the relevant titration data and results. You should calculate the molarity
for each of the three good fine titration trials you did, then average those. Make a neat, well-labeled
copy of your scale-up calculation and one sample calculation of your unknown’s molarity. Be sure to
write your name on your report and also record your unknown vial number and the volume of diluted
bleach solution you used for your fine titrations near your data table. Staple all scratch paper to the back
of this. Your instructor will initially grade only your average molarity of NaOCl, but if it is “way off”
they might let you redo the calculations for a penalty.
66
Experiment 11: Redox Reactions & Equations
For this experiment there is a station set up for each reaction. Instead of working in your typical
groups of two you should pair up into groups of four. It does not matter what order you do the stations,
but some will be demonstrations that your instructor will do for you first. Carefully write down your
own observations (each person) for each reaction; be sure to write down reactant descriptions,
immediate changes observed when combined, and changes observed after some time has passed . After
your group has completed the reactions you need to go through the process shown in the redox example
below to determine the redox reaction for the process. You will need to write and balance your own net
ionic equation for each reaction. As shown in the example you should look for half-reactions involving
the chemicals you mix to get started on the equation writing (use the table of standard reduction
potentials at the end of this experiment). Your instructor will give you hints if you need some. Many of
the equations will be relatively simple and not involve that many different species—look for simple
equations first! Once you determine the reaction equation answer any additional questions concerning
the process (this applies for reaction #1 and #3).
You can work with others as you complete these tasks, but ultimately you need to come up with and
write out your own answers. You will be penalized if your caught simply copying answers from others;
in a previous semester a number of students received zeros on this lab for sharing equations. If
discussing a process with other students does not help you come up with the answer ask your instructor,
DON’T copy answers from other students.
Redox Example
The reaction of manganese(II) chloride with sodium bismuthate (NaBiO 3 ) in acidic media.
Lab observations:
Manganese (II) chloride tetrahydrate is a pale pink solid. When it is mixed in water, a colorless
solution results. Addition of the yellow sodium bismuthate with stirring results in a pale purple solution
with no solid present.
What is the redox reaction for this process?
•
•
What is present?
o 𝑀𝑛2+ , 𝐶𝑙 − , 𝑁𝑎+ , 𝐵𝑖𝑂3−
Look at the available half reactions (use the table at the end of this experiment) and find half
reactions that contain your beginning components. Suppose the following are what we find.
o Available half-reactions with 𝑀𝑛2+
▪ 𝑀𝑛𝑂4− (𝑎𝑞) + 8 𝐻+ (𝑎𝑞) + 5 𝑒 − → 𝑀𝑛2+ (𝑎𝑞) + 4 𝐻2 𝑂(𝑙 )
▪ 𝑀𝑛𝑂2 (𝑠) + 4 𝐻 + (𝑎𝑞) + 2 𝑒 − → 𝑀𝑛2+ (𝑎𝑞) + 2 𝐻2 𝑂(𝑙 )
▪ 𝑀𝑛2+ (𝑎𝑞) + 2 𝑒 − → 𝑀𝑛(𝑠)
1.51 V
1.23 V
-1.17 V
o Available half-reactions with 𝐶𝑙 −
▪ 𝐶𝑙 2 (𝑔) + 2 𝑒 − → 2 𝐶𝑙 − (𝑎𝑞)
1.36 V
o Available half-reactions with 𝑁𝑎+
▪ 𝑁𝑎+ (𝑎𝑞) + 𝑒 − → 𝑁𝑎(𝑠)
-2.71 V
o Available half-reactions with BiO 3 ▪ 𝐵𝑖𝑂3− (𝑎𝑞) + 2 𝑒 − + 6 𝐻+ (𝑎𝑞) → 𝐵𝑖 3+ (𝑎𝑞) + 3 𝐻2 𝑂(𝑙 )
2.03 V
67
•
Find the best oxidizing agent and best reducting agent, consistent with experimental
observations.
o If we blindly choose the most positive and most negative voltages, we’d be using
▪ 𝐵𝑖𝑂3− (𝑎𝑞) + 2 𝑒 − + 6 𝐻+ (𝑎𝑞) → 𝐵𝑖 3+ (𝑎𝑞) + 3 𝐻2 𝑂(𝑙)
2.03 V
+(
−
)
(
)
▪ 𝑁𝑎 𝑎𝑞 + 𝑒 → 𝑁𝑎 𝑠
-2.71 V
o We need one oxidation and one reduction, so we need to reverse the most negative one:
▪ 𝐵𝑖𝑂3− (𝑎𝑞) + 2 𝑒 − + 6 𝐻+ (𝑎𝑞) → 𝐵𝑖 3+ (𝑎𝑞) + 3 𝐻2 𝑂(𝑙)
2.03 V
+
−
▪ 𝑁𝑎(𝑠) → 𝑁𝑎 (𝑎𝑞) + 𝑒
2.71 V
o Adding these two together gives Na(s) as a reactant, and we have no Na (s) present, so
we can’t use that half-reaction.
o The next-most negative voltage is the last listed reaction with manganese, but reversing
it would give Mn(s) as a reactant, and that is not present.
o The third-most-negative one is
▪ 𝑀𝑛𝑂2 (𝑠) + 4 𝐻 + (𝑎𝑞) + 2 𝑒 − → 𝑀𝑛2+ (𝑎𝑞) + 2 𝐻2 𝑂(𝑙 )
1.23 V
o Reversing it results in reactants that are present, but it also shows the formation of solid
product, which did not occur.
o The fourth-most-negative choice would be
▪ 𝑀𝑛𝑂4− (𝑎𝑞) + 8 𝐻+ (𝑎𝑞) + 5 𝑒 − → 𝑀𝑛2+ (𝑎𝑞) + 4 𝐻2 𝑂(𝑙 )
1.51 V
o Reversing it gives
▪ 𝑀𝑛2+ (𝑎𝑞) + 4 𝐻2 𝑂(𝑙 ) → 𝑀𝑛𝑂4− (𝑎𝑞) + 8 𝐻+ (𝑎𝑞) + 5 𝑒 −
-1.51 V
2+
o 𝑀𝑛 and water are reactants, (and those are present) and it gives only soluble
products, which matches our observations in lab.
o Thus, if we take the best voltages consistent with what is present and what is observed,
we would use:
▪ Oxidation:
• 4 𝐻2 𝑂(𝑙 ) + 𝑀𝑛2+ (𝑎𝑞) → 𝑀𝑛𝑂4− (𝑎𝑞) + 8 𝐻+ (𝑎𝑞) + 5 𝑒 −
▪ Reduction:
• 2 𝑒 − + 6 𝐻 + (𝑎𝑞) + 𝑁𝑎𝐵𝑖𝑂3 (𝑠) → 𝐵𝑖 3+ (𝑎𝑞) + 3 𝐻2 𝑂(𝑙 ) + 𝑁𝑎 + (𝑎𝑞)
o The reaction is balanced by scaling the two half-cell reactions to involve the same
number of electrons (multiplying the oxidation reaction by the number of electrons in
the reduction step and vice versa):
▪ 8 𝐻2 𝑂(𝑙 ) + 2 𝑀𝑛2+ (𝑎𝑞) → 2 𝑀𝑛𝑂4− (𝑎𝑞) + 16 𝐻 + (𝑎𝑞) + 10 𝑒 −
▪ 10 𝑒 − + 30 𝐻+ (𝑎𝑞) + 5 𝑁𝑎𝐵𝑖𝑂3 (𝑠) → 5 𝐵𝑖 3+ (𝑎𝑞) + 15 𝐻2 𝑂(𝑙 ) + 5 𝑁𝑎+ (𝑎𝑞)
o Adding these two reactions eliminates the electrons terms and yields the balanced
reaction:
▪ 14 𝐻+ (𝑎𝑞) + 2 𝑀𝑛2+ (𝑎𝑞) + 5 𝑁𝑎𝐵𝑖𝑂3 (𝑠)
→ 7 𝐻2 𝑂(𝑙 ) + 2 𝑀𝑛𝑂4− (𝑎𝑞) + 5 𝐵𝑖 3+ (𝑎𝑞) + 5 𝑁𝑎+ (𝑎𝑞)
68
Name: ______________________________
Procedure
1. a) Penny plus HCl
b) Penny plus nitric acid
Note: You need to prepare the HCl part early in the lab period (at least within the first hour) and
leave it until close to the end of the lab period to take your final observations (it will not be done
reacting, but it will be far enough to understand what is happening).
Bring at least two pennies to lab; they need to be post 1983. Put your name on a label and place it
on a small beaker. Use a triangular file to scrape the edge of a post 1983 penny (it seems to work best
to scrape in multiple places on opposite edges of the penny) or cut the penny in half if you wish (and
then use both halves).
Part a. Pour about 50 mL of 6 M hydrochloric acid into a small beaker. Gently slide the penny
down the side of the beaker into the hydrochloric acid; don’t have your face close when you do this,
in case it splashes! Observe for a few minutes, then place the beaker at the back of a hood and put
a watch glass on top of it to keep dust out. Make sure there is a sign on this hood that says “Keep
heat and flames away from this hood”. Come back and take your final observations before the end
of the lab period.
Part b. Repeat part a but with HNO 3 instead of HCl. This one will react quicker than the HCl
so you don’t have to wait as long to take your final observations.
Disposal: Pour the liquids down the drain and flush with lots of water. Do not lose what remains of the
penny in the drain! It is easiest to pour the beaker contents through a funnel that will catch what remains
of the penny and then rinse it so you don’t get remaining acid on your hands. The remains of the penny
can be disposed of in the trash.
Question: What caused the difference in product in (a) and (b)—consider Eo ’s of reactions in your
answer. What does the difference in (a) and (b) tell you about the strength of the HCl and HNO 3 as
oxidizing agents? Why did I specify a post 1983 penny? What would happen you used a pre-1980
penny in the HCl? In the HNO 3 ?
2. Al and CuCl2
Place about a pea size amount of CuCl2 in a vial and add just enough water to dissolve it (stir with
stirring rod to get solid to dissolve). Lightly crumple a piece of Al foil or powdered Al (whichever is
available) so it will go down into the vial. Push the foil down into the solution of CuCl2 with the stirring
rod and observe carefully for a few minutes, then check it every few minutes until it quits changing
appearance.
Note: A gas is sometimes evolved when these are mixed. This is a side-reaction and can be ignored.
Keep this equation simple.
Disposal: Remove any solid left (pour through filter paper if necessary), rinse it with water and put it
in the trash can. The solution leftover can go down the sink (run the water several minutes afterward).
69
3. “Elephant toothpaste”
CAUTION--30% HYDROGEN PEROXIDE BURNS SKIN QUICKLY! DON’T GET IT ON YOU.
WASH UP ALL SPILLS IMMEDIATELY (THIS INCLUDES ANY THAT DRIPS DOWN THE
SIDES OF THE BOTTLE.)
While wearing gloves, pour about 5 mL of 30% hydrogen peroxide into a 50 mL graduated cylinder.
Observe it carefully. Bubbles indicate the H 2 O 2 is decomposing. Add 2 mL of dish detergent so any
bubbles forming will result in foam; this is just so the bubbles will be easier to see, and otherwise does
not influence the reaction. After observing again, place the graduated cylinder into the sink and add a
pea size amount of KI, which is a catalyst. A catalyst speeds up a reaction but is not itself consumed;
you may see a brown/orange color from the catalyst—you don’t need to explain this with your reaction
equation.
Hint for the reaction equation: This is a decomposition reaction, so the only reactant in the final
equation is the hydrogen peroxide. However, each half-reaction can have other stuff that could cancel.
Look for one half-reaction in which hydrogen peroxide is oxidized and another in which it is reduced.
Disposal: When the reaction is over, be sure to clean up glassware (inside and outside) while wearing
gloves. Make sure to rinse the outside of the gloves well and leave them for the next person. All products
from this reaction can be rinsed down the sink. Run water for a few minutes after putting them in the
sink.
Questions: If the hydrogen peroxide is decomposing, what might be forming? Come up with several
possible reactions that would explain your observations and use half-cell potentials to determine which
is most likely. Note that you don’t need to explain the catalysis.
4. Zn in copper(II) sulfate
Add a small piece of zinc to a solution of copper(II) sulfate in a vial. Don’t use much of the solution
as you won’t need it to see the reaction. Check it every five minutes. Keep this equation simple.
Disposal: Use tongs or forceps to fish the wire out of the vial, holding it over the sink while you do so.
Rinse the wire off with water from a squeeze water bottle. Place the solution from the vial in heavy
metal waste container. Rinse the sink very thoroughly, even if you don’t think you got anything in it.
5. Sodium and water (demo).
Your instructor will do this as a demonstration for you. He/she will first place a small piece of
sodium metal into the water. After that reaction has stopped, he/she will add some phenolphthalein
indicator to help you figure out the products of the reaction. This equation is not too bad when you take
all the evidence into account.
6. Iron(II) Ion and Permanganate Ion
Dissolve a small (~ ¼ of a pea) sized amount of iron(II) ammonium sulfate in a minimum amount
of water. Add 2-3 drops of the provided sulfuric acid and then add the provided KMnO 4 solution
dropwise.
70
7. Thermite (demo)
Your instructor will do this as a demonstration for you. You should observe the Al and Fe 2O3
reactants before they are mixed as well as what occurs during the reaction and what the products look
like. One product will be molten elemental iron. There is no gaseous product. This demonstration will
be done outside, so you will all have to do it at the same time. Keep this equation simple, but remember,
solids do not form of just one charge of ion.
Some Standard Reduction Potentials at 25 oC*
Half-Reaction
E° (V)
K+(aq) + e- -----> K(s)
-2.93
Na+(aq)
-2.71
+
e-
-----> Na(s)
Mg2+(aq) + 2 e- -----> Mg(s)
-2.37
Al3+(aq)
-1.66
+3
e-
-----> Al(s)
Mn2+(aq) + 2 e- -----> Mn(s)
-1.18
2 H 2 O + 2 e- -----> H2(g) + 2 OH -(aq)
Zn2+(aq) + 2 e- -----> Zn(s)
-0.83
-0.76
Fe2+(aq) + 2 e- -----> Fe(s)
-0.44
Ni2+(aq)
e-
+ 2 -----> Ni(s)
Sn2+(aq) + 2 e- -----> Sn(s)
-0.44
-0.23
Pb2+(aq) + 2 e- -----> Pb(s)
-0.13
2 H +(aq) + 2 e- -----> H2(g)
0.00
Sn4+(aq)
+0.13
+2
e-
----->
Sn2+(aq)
Cu2+(aq) + e- -----> Cu+(aq)
+0.13
SO 4 2-(aq) + 4 H+(aq) + 2 e- -----> SO 2(g) + 2 H2 O
+0.20
Cu2+(aq)
+ 2 -----> Cu(s)
I 2(s) + 2 e- -----> 2 I -(aq)
+0.34
+0.53
MnO4 -(aq) + 2 H 2 O + 3 e- -----> MnO2(s) + 4 OH -(aq)
+0.59
O2(g) + 2 H +(aq) + 2 e- -----> H2 O2(aq)
Fe3+(aq) + e- -----> Fe2+(aq)
+0.68
+0.77
NO 3 -(aq) + 4 H+(aq) + 3 e- -----> NO (g) + 2 H2 O
+0.96
Br2(l) + 2 e- -----> 2 Br-(aq)
+1.07
H +(aq)
+1.23
O2(g) + 4
e-
+4
e-
-----> 2 H2 O
MnO2(s) + 4 H+(aq) + 2 e- -----> Mn2+(aq) + 2 H 2 O
Cl2(g) + 2
e-
-----> 2
Cl-
+1.23
+1.36
(aq)
MnO4 -(aq) + 8 H +(aq) + 5 e- -----> Mn2+(aq) + 4 H2 O
H2 O2(aq) + 2 H+(aq) + 2 e- -----> 2 H2 O
+1.51
+1.77
O3(g) + 2 H +(aq) + 2 e- -----> O2(g) + H2 O
+2.07
71
Exam and Checkout
In the first hour of the last lab period, you will take an exam which is covers material from both
CHEM 111 and CHEM 112. This is a standardized ACS exam that will count as one lab grade.
Additionally, it will give you practice for the final exam in lecture which will also be a standardized
ACS exam, but only over CHEM 112 material.
In the second part of the lab period, you will check and clean the equipment in your drawer and do
an additional lab cleaning job assigned by your instructor.
Failure to attend Lab Check Out and Exam will result in the failure of both 112
Lab and Lecture.
72