1

Lab #
1
Date
7/3
7/8
Experiment
Introduction to course, check-in, Molar Mass from
Freezing Point Depression
Calcium Iodate (R)*
Page
8
2
3 7/10 Rate and Activation Energy of Iodination of Acetone
10
12
16 4
5
7/15
7/17
Iron (III) Nitrate and Potassium Thiocyanate (R)
Acetic Acid in Water
Acid Base Titrations and Indicators, and Buffers
18
6 7/22 20
7 7/24
8 7/29 Redox Reactions and Electrochemical Cells, Checkout 25
*Experiments with an (R) will require formal laboratory reports. These reports will be due at 5:00
PM the Wednesday after the experiment is completed.
2

Chemistry 210L – General Chemistry Lab II, Spring 2014
Instructor Contact Information:
Section:
-
Instructor: Office: Phone: Email
Asoka Marasinghe HA 407E 477-2277 asoka@mnstate.edu
Required Materials: Laboratory notebook with carbon-copy pages (MSUM bookstore)
Safety Goggles
CHEM 210L General Chemistry Laboratory II (1) Laboratory techniques of general chemistry including qualitative and quantitative analysis. Course should be taken concurrently with CHEM
210. Prerequisite: CHEM 150
Desire-2-Learn (D2L): https://mnstate.ims.mnscu.edu/
D2L e-mail system will be used to distribute information for lab. Email messages sent from D2L are sent to your mnstate.edu email account. Be sure to check your mnstate.edu email account regularly.
Lab Manual:
The experimental procedures for this class (lab manual) will be posted on my website at http://web.mnstate.edu/marasing/CHEM210L_Summer/Chem210L.htm It is highly recommended that you print the entire manual out (preferably two-sided as it is set up for this type of printing) at the beginning of the semester and bring it with you to each lab period.
Pre-lab Videos, Prelab meeting and Quizzes: 5 pts each quiz
There will be must see video for each lab. The videos are hosted on YouTube and can also be accessed directly at http://www.youtube.com/user/drbodwin . There are links to them on my website.
There will also be a required pre-lab quiz for each lab in D2L, due before the start of a lab (see D2L for deadline). Late quizzes will not be accepted for any reason. Penalties for missed quizzes are: first missed quiz = zero points for the quiz, regular points for the rest of the experiment; subsequent missed quizzes = zero points for the quiz, half credit for the rest of the experiment. A “missed” quiz is any quiz on which 0 points are earned. There will be a prelab meeting in SL102 at the beginning of each lab, do not miss it.
Arrival and Attendance: Late arrival penalized on Notebook Copy points
Late arrival in lab is very disruptive and should be avoided. If you know you will be late for a lab, inform the instructor and your lab partner before class. Penalties for late arrival: first late arrival is excused; second late arrival is half credit on Notebook Copy points; subsequent late arrivals result in no Notebook Copy points. Anyone more than 10 minutes late (as determined by the instructor) will not be allowed to participate in the experiment and will receive a grade of zero for the experiment.
Make-up labs will not be permitted without advance notice and approval by the instructor.
Notebook Copies: 5 pts each week
You are required to turn in the carbonless copies of your lab notebook for each lab excercise . All experimental deviations from the lab manual, observations made during the experiment, data collected, and analysis should be recorded in your notebook as the experiment is being performed. If the notebook copy pages are not turned in at the end of the lab period, no points will be earned for the hand-in assignment or lab report associated with the experiment.
3

For data collected or analyzed by computer, you must still record information in your lab notebook, but not all data needs to be transcribed or reproduced completely in your notebook. For data generated in LoggerPro, briefly describe the data in your lab notebook, for example “The graph of absorbance vs. wavelength showed a single peak centered at wavelength 545nm” or “The temperature increased linearly from 17.4°C to 35.8°C over 3.68 minutes”. For calculations performed by the computer (LoggerPro, Excel, etc.), describe and show a sample calculation for each formula or type of calculation you did. For example:
The moles of gallium used in each experiment was calculated by dividing the mass of gallium pellets by the atomic mass of gallium,
69.723
g / mol
. Using values from experiment 1:
6.831
g
Ga
´
1  mol   Ga
69.723  g
Ga
=
0.09797  mol
Ga
If you have saved the information in a file, it’s not a bad idea to note the location and filename you used so you can find the data more easily when you are working on the assignment.
Hand-In Assignments (HI) and Lab Reports (LR): 20 pts (HI) or 30 pts (LR) each
Each experiment will have a hand-in assignment or a lab report. The assignments of Tuesday labs due by 5:00 PM on the following Friday (following completion of the experiment in the lab) and for
Thursday labs are due by 5:00 PM on the following Monday. Assignments more than one week late will receive zero points.
All assignments must be typed and submitted to the labeled slot of the drop-box located in front of
Hagen 103 before the submission deadline . All associated graphs and calculations must be saved in an electronic format as part of the Microsoft Word file in order to enable them to be easily read.
The blank hand-in assignment will be posted as a Microsoft Word document in my website and will include a grading rubric that may be followed to make sure that you have completed everything required for the hand-in.
Safe Practices: 10 pts
Everyone is expected to follow safe and considerate procedures in lab. At the beginning of the semester, everyone starts with 10 points which are lost if unsafe practices are observed. This includes not wearing goggles, not cleaning up spills, not replacing covers on stock bottles, etc.
Except in extreme cases, a warning will typically be given before points are lost.
Grading:
There will be 8 experiments and one report-writing workshop (8 pre-lab quizzes and 8 hand-ins or lab report assignments) over 5 lab weeks (8 notebook copies). There is a tentative total of approximately 280 points for the course, although this may change slightly. Tentative grade ranges:
A = 100-90%, B = 89-80%, C = 79-70%, D = 69-60%. These cutoffs may be lowered at the instructor’s discretion, but will not be raised.
Academic Honesty: { http://www.mnstate.edu/sthandbook/POLICY/index.htm
}
Cheating will not be tolerated. You will be working with a partner in the lab, and you are welcome to work together on data analysis, but each student will be required to submit individual assignments that are not duplicates of your partner or anyone else.
The minimum penalty for copied work will be a grade of zero for the assignment but academic dishonesty may lead to expulsion from the University. You are responsible for protecting the integrity of your own work so
DO NOT send your partially or fully completed assignments to anyone. If they turn it in as their own
4

you will both suffer the consequences of any disciplinary action whether you approved of the copying or not. For a full description of the MSUM Academic Honesty Policy, see the Student
Handbook.
Disability Access Statement : Students with disabilities who believe they may need an accommodation in this class are encouraged to contact Greg Toutges, Director of Disability Services at 477-4318 (Voice) or 1-800-627-3529 (MRS/TTY), Flora Frick 154 as soon as possible to ensure that accommodations are implemented in a timely fashion. Information regarding Disability Services is available at http://web.mnstate.edu/disability/
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Features of a Sample Report Format
Title of Report
Author’s Name
Partner’s Name(s)
Date of Experiment
Abstract:
Provide a brief description of what you determined in this experiment. Obviously, you won’t be able to write this section until you have completed the experiment and have a good grasp of the concept that was developed as a result of interpreting the data.
Usually there is (1) a statement of the purpose/problem, followed by a (2) very brief statement of method, a (3) brief statement of results including numeric values determined where appropriate, and finally the (4) conclusion or implications. Write the abstract LAST .
Experimental:
[Draw a labeled diagram of the experimental set up if necessary . Don’t draw trivial items.]
Give a description of the actual technique used in the experiment. As appropriate, use a labeled diagram to describe the experimental setup. This section should be written in past tense and passive voice. For example, “Hydrochloric acid was added to the solution ...” . NOT simple past tense “Next I added HCl to the solution” and NOT a simple command “Add HCl to the solution”. Indicate the equipment used and all quantities you measured (refer to tables as necessary for quantities). Do not duplicate the procedure as written in your lab manual; summarize or paraphrase. Avoid stating procedures you can assume the reader of the report knows
, e.g., “The spectrophotometer was calibrated by first turning on the power switch...” Mention should be made either here or in the results/discussion of the purpose of steps in the procedure that are not self evident. For instance, the nesting of two beakers inside each other is a critical step in one lab. It helps to reduce a certain kind of error, and by stating the reasoning for these steps, you will achieve a higher grade.
Results and Discussion:
This section will generally include all observations, data and calculations involving data
(including graphs) obtained from the laboratory and answers to any questions in the laboratory manual . All graphs should be clearly labeled with proper axis labels (and units), names of authors, etc. All data and calculations should be presented in an organized fashion using Equation
Editor. Tabulate (put in an organized table, rather than listing on a single line) data whenever possible. When generating graphs, make sure that the graph and the associated legends etc. are large enough to clearly show the data and any trends. For some experiments, a relatively small graph embedded in the text is sufficient; for other experiment, the amount and complexity of data presented in the graph make it difficult to present clearly unless the graph is on a separate page. If a graph is on a separate page, it should cover at least 75% of the page. All data and results should indicate proper significant figures and appropriate calculated errors, if appropriate. Sources of error (often a result of a limitation in the procedure) should also be included in this section. This shows a reader the limitations of the techniques employed and identifies areas that could be improved in the future.
The questions in the laboratory manual should be addressed in this section in narrative form (not
“Yes, No, Blue, 10 ˚C” but rather “as the central atom of the polyatomic ion moved to the right on the periodic table, an overall trend of increasing pH was noted, consistent with notes presented in lecture”), referring back to the experimental results you have obtained. The question number (in parentheses before the discussion statement) may be indicated for clarity. These questions are designed to guide you through the important features of the laboratory exercise and to demonstrate to the instructor that you have understood the principles involved.
6

Conclusions:
This section should reflect the overall results of the experiment. What does it all mean? Ask yourself, what did I learn; what is the final/overall result from the experiment? This section should be short! If you find yourself writing extensive amounts in this section they should probably be written in the discussion section! Some summarized data should be presented, in many cases it will be the same bit of data that will be included in the abstract.
7

When a pure substance freezes, the individual particles (molecules, atoms, ions) must organize themselves such that their intermolecular forces prevent them from flowing. If an impurity is introduced into this pure substance, the intermolecular forces that allow the substance to freeze are disrupted meaning that more energy must be removed from the system before the particles can organize. This means that a solution will always freeze at a lower temperature than the pure solvent.
It is important to note that the amount that the freezing point is lowered does not depend upon the identity of the solute; it’s just dependent upon the number of solute particles in the solution.
Properties which depend only upon the amount of solute and not its identity are called colligative properties, and they include freezing point depression, boiling point elevation and others. The concentration unit often used to calculate colligative properties is molality because the molality of a solution is independent of the temperature. Molality has units similar to molarity, but rather than mols of solute /
L of solution
(molarity, M), molality is defined as mols of solute / kg of solvent
and is abbreviated with a lower-case “m”.
Cyclohexane is extremely flammable. Do not have any flames in the laboratory.
I. Freezing Point of Pure Cyclohexane
Obtain an 18x100mm test tube and place a small Teflon-coated stir bar in the bottom of the tube.
Pipette 5.00 mL (±0.01 mL) of cyclohexane into the test tube. Clamp the tube on the ring stand inside a 400-mL beaker. Add a stir bar and clamp the temperature probe in the tube. The probe should not touch the wall of the tube and should be centered in the test tube as much as possible.
Once the test tube and temperature probe are positioned, slide the stir plate out from beneath the apparatus and remove the beaker. Prepare an ice-water bath in the beaker (mostly ice, just enough water to make the bottom half or two-thirds look wet). Set up LoggerPro to record temperature vs. time. When you are ready to begin a run, click Start and then slide the ice-water bath onto the test tube and position the stir plate. Collect temperature vs. time data until the sample is frozen and record the freezing point. Remove the ice-water bath, melt the cyclohexane sample and repeat until you are confident in your results.
Save one representative cooling curve for cyclohexane on which you have indicated the freezing point.
Be sure that each person saves a properly formatted graph. The easiest way to attach the graph to your hand-in is to use the Snipping Tool accessory and paste the graph directly into a
Microsoft Word file. Find the Snipping Tool accessory under the Start menu: All programs:
Accessories: Snipping Tool. Select only the graph itself, not the data, when using the Snipping
Tool. Paste the graph directly into an empty Microsoft Word file and email this file to yourself.
Then you can easily copy and paste the graph into your hand-in at a later time.
Q: Why should the temperature probe NOT be allowed to touch the side of the test tube? Why are you measuring the freezing point of cyclohexane and not simply looking it up?
Compare your freezing point to others in the class and to published/accepted values.
II. Freezing Point Depression of a Cyclohexane Solution
A. Weigh approximately 0.080-0.090g of benzophenone (C
6
H
5
COC
6
H
5
) to the nearest thousandth of a gram (±0.001g) and record the mass. Remove the temperature probe from the test tube being careful not to lose any liquid from the tube and add the sample of
8

benzophenone to the cyclohexane. Allow the benzophenone to dissolve by stirring and warming the solution if necessary. Reposition the probe and perform the same procedure as in part I to determine the freezing/melting point of the resulting solution. Each student should save one representative cooling curve for this solution to a Microsoft Word file just as you did in part I. Be sure that the freezing point is noted on the graph.
Q: Did the freezing point of the solution differ from the freezing point of the pure solvent?
Explain your observation on a molecular level.
B. Calculate the freezing point depression constant, k fpd
, for cyclohexane. (ΔT fp = k fpd
·m·i)
Benzophenone is a molecular solute in cyclohexane; therefore, each benzophenone molecule yields one solute particle so the value of “i” is 1.
III. Determination of the Molar Mass of an Unknown
A. Empty the cyclohexane solution from part II into the appropriately labeled waste container in the hood. Rinse the tube with several 1-2 mL portions of solvent from a wash bottle. Dry the tube and clean the magnetic stir bar and the probe. Dispense a fresh 5.00-mL aliquot of cyclohexane into the tube. Bottles containing unknowns will be placed by the balances with specified ranges of masses to be used for each unknown. Weigh out (to ±0.001 g) the appropriate amount of the unknown assigned to you. Be sure to record your unknown number in your lab notebook and to communicate it in your lab report or hand-in. Please leave the balance pan and the balance area clean for other students and for yourself. Measure temperature vs. time data as above to determine the freezing point of this solution. Each student should save one representative cooling curve for this solution to a Microsoft Word file as you did in the previous sections.
Q: How can you determine the concentration (molality) of this solution if you do not know the identity of the solute? {NOTE: You may assume that all the unknowns are molecular solutes.}
Each part of this experiment uses the same equation; the only thing that changes is the unknown variable. In the first part, we need to determine the freezing point of pure cyclohexane so we can determine how much the freezing point temperature changes when a solute is added. In the second part, we are using a known solute so we can make a solution of known concentration. With that solution of known concentration, we can measure the change in freezing point temperature and use this data to calculate the freezing point depression constant (k fpd ) for cyclohexane. Once we know the value of the freezing point depression constant for cyclohexane, we can use it to determine the molar mass of an unknown molecular solute.
(T fp,puresolvent
-
T fp,solution
)
= D
T fp
= k fpd
× m
× i
=
( )
×
× grams of  solute
×
×
× molar mass of  solute
× kg of  cyclohexane solvent
×
×
×
×
( )
Part I: Determine the freezing point of pure cyclohexane
Part II: Determine the freezing point depression constant for pure cyclohexane
9
Part III: Determine the molar mass of an unknown molecular solute

A common method for purification and isolation of various substances involves the precipitation of either the desired substance or impurities from an impure mixture. For this method to work, it is desirable that all of the desired precipitate forms and can be removed from the mixture, but this is not typically the case. It is much more common that some of the “precipitate” remains dissolved in the solvent, meaning that pure compounds can be difficult to obtain in high yield in this manner.
There are no specific concerns for this experiment. Wear goggles at all times and dispose of all waste in the appropriate container.
I. Precipitation Reaction between Calcium Nitrate and Potassium Iodate
Using appropriate graduated cylinders, prepare the following mixtures of Ca(NO 3 ) 2 and KIO 3 in clean labeled 150 mL beakers.
Solution #1 25.0 mL Ca(NO 3 ) 2 + 25.0 mL KIO 3
Solution #2 15.0 mL H 2 O + 10.0 mL Ca(NO 3 ) 2 + 25.0 mL KIO 3
Mix each solution with a separate stirring rod. Do not remove the stirring rods from the solution or you may lose some of the solid that should be recovered and weighed later on. If solid is not forming, you may try scratching the sides and bottom of the beaker with the stirring rod to help induce crystallization. Let these solutions stand for at least 15 minutes. Stir occasionally.
The following questions are designed to help you interpret your data. It may be helpful for you to answer these questions while you are in the lab and can interact with the instructor but you are not required to record the answers to these questions in your lab notebook.
Q: In what ways are these two solutions different? What role does the water play in the second solution? Determine the limiting reactant and theoretical yield of solid for each of the solutions.
II. Experimental Mass of the Precipitates
Label two filter paper circles with pencil. Weigh the filter papers. Following the directions given by your lab instructor, filter each solution. When the filtration has been accomplished, set the filtrates aside for use in part III. DO NOT THROW AWAY YOUR FILTRATES.
Rinse the precipitates with methanol in the fume hood to help remove residual liquid from the solids.
Remove the filter papers from the funnels and place them on watch glasses. Dry in air for several minutes and then place them in a 90-110ºC drying oven for at least 45 minutes. Allow the samples to cool to room temperature before recording their final mass.
Perform part III while your samples are drying in the oven.
Q: What was your percent yield for each sample? Why was it not 100%?
III. Qualitative testing of the filtrate
A. Standards - Characteristic tests for the presence of Ca
2+
and IO
3
–
. The purpose of these tests is to see what happens with (relatively) high concentration of the test ion, relatively low concentrations of the test ion, and with no test ion present. Make careful observations so you can correctly interpret the results of these tests on your filtrates in Part B.
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1. Qualitative Tests for Ca
2+
3 Ca
2+
(aq) + 2 PO
4
–3
(aq)

Ca
3
(PO
4
)
2
(s) a. Obtain about 1 mL of Ca(NO drops) of Na 3 PO 4
3
)
2
solution in a small test tube. Add about 0.5 mL (10
solution. This is a strong positive test for Ca
2+
(aq). b. Obtain 2-3 drops of Ca(NO
3
)
2 mL (10 drops) of Na
3
PO
4
solution and add about 1 mL of water. Add about 0.5
solution. This is a weak positive test for Ca
2+
(aq). c. Repeat this test with 1mL of water, no added Ca(NO 3 ) 2 solution.
Q: Are your observations different? Why?
2. Qualitative Tests for IO
3
IO 3
–
–
(aq) + 5 I
–
(aq) + 6 H
+
(aq)

3 I 2 (aq) + H 2 O(l) a. Obtain about 1 mL of KIO 3 solution in a small test tube. Dissolve a small amount of
KI (about the size that would fit on a match head) in the solution and then add about 1 mL of 1 M HCl. This is a strong positive test for IO 3
-1
(aq). b. Obtain 1-2 drops of KIO
3
solution in a small test tube and add about 1 mL of water.
Dissolve a small amount of KI (about the size that would fit on a match head) in the solution and then add about 1 mL of 1 M HCl. This is a weak positive test for
IO
3
–
(aq). c. Repeat this test with 1mL of water, no added KIO 3 solution.
Q: Are your observations different? Why?
B. Filtrates (Note that in this section you will perform a total of 4 tests, 2 on each filtrate.)
1. Test both your filtrates for the presence of Ca
2+
. Perform the same test for Ca
2+
as above
(1a), but instead of the known Ca
2+
solution, use about 1 mL of filtrate.
Q: Was calcium present in your filtrates? If so, based upon your observations, how much
Ca
2+
was present (a significant amount or just a little bit)? Are these observations consistent with the identity of the limiting reagent in each solution?
2. Test both your filtrates for the presence of IO
3
(2a), but instead of the known IO 3
–
–
. Perform the same test for IO
solution, use about 1 mL of filtrate.
3
–
as above
Q: Was iodate present in your filtrates? If so, based upon your observations, how much IO
3
– was present (a significant amount or just a little bit)? Are these observations consistent with the identity of the limiting reagent in each solution?
Review all sections of this experiment. We often describe salts as being either “soluble” or
“insoluble”. For example, you might say “Sodium chloride is soluble” or “Calcium carbonate is insoluble.” Given your observations throughout this experiment, why do these statements not completely describe the concept of “solubility”?
11

The rate law and the activation energy for this reaction will be determined by measuring the initial rates of the reaction while systematically varying the concentration of reagents and the temperature of the reaction of acetone with iodine.
O
H
+
O
C
+ I
2 C
+ I
-
+ H
+
H
3
C CH
3
H
3
C CH
2
I
The above reaction is catalyzed by hydrogen ions, H
+
(i.e. H 3 O
+
), so this ion appears in the rate law, which takes the form: rate = k[acetone] a
[H
+
] b
[I
2
] c where a, b, and c are the reaction orders with respect to acetone, H
+
and I 2 respectively. This is a convenient reaction to study because the reaction order with respect to iodine is zero (c=0), so the reaction rate does not depend on the concentration of I 2 . The reaction will proceed as long as iodine is present. Also note that I 2 a characteristic yellow-brown color (especially in the presence of I
–
, which forms I 3
–
in solution). The
has experiment is designed so that a large excess of acetone and hydrogen ions is employed and a much smaller but known concentration of iodine is used. Using the following equation, the rate of the reaction can be calculated. rate
 
Δ[I
2
]
Δt
CAUTION: Iodine will stain balances, skin, and clothing. If a small spill occurs, rinse with a solution of
Na 2 S 2 O 3 .
I. Stock Solutions
Obtain the following stock solutions, recording the exact, actual concentrations listed on the stock bottles:
20 mL of Acetone(aq) (approx. 4M)
20 mL iodine solution (approx. 5mM)
20 mL HCl(aq) (approx. 1M)
II. Preparation of Reaction Mixtures and Data Collection
A. Obtain ~250 mL of room temperature water in a 400 mL beaker and place it on a white sheet of paper to give a bright white background for observing the color of the solutions in the reaction tubes. In all the following runs use pipettes to dispense the solutions.
B. For Run #1 prepare the solutions as follows. Add 2.00 mL each of ~4 M acetone(aq), ~1 M H
+
(aq) and distilled/deionized water to a large test tube (labeled 1A), and swirl to mix these reagents well. To a second large test tube (labeled 1B), add 2.00 mL of I 2 solution. Place tubes 1A and 1B in the water bath to allow their temperatures to equilibrate.
12

C. Data Collection: Run #1.
Open LoggerPro on the computer. The program should detect the SpectroVis Pro and set up a default data collection. If it does not, ask your instructor or Lab Assistant for help.
1. Collect an Absorbance vs. Wavelength Spectrum of Iodine a. Prepare a sample by diluting a few drops of the iodine stock solution to 2-3mL with water, enough to fill a cuvette. This does not have to be an exact dilution; it should be dark enough to produce a strong spectrum but not so dark that the maximum absorbance is above ~1. b. Collect a full spectrum and print it, noting the region(s) of high absorbance.
Q. Is the spectrum you collected consistent with the observed color of the solution? For this part of the experiment, why is it not necessary to know the exact concentration of the iodine sample you are measuring? How will your Absorbance vs. Wavelength spectrum change if you use a less dilute sample? A more dilute sample?
2. Setting Up LoggerPro to Measure Reaction Times a. From the “Experiment” menu in LoggerPro, choose “Data Collection” and switch to a “Time
Based” experiment with the following parameters:

Experiment length = 10 minutes
 Continuous sampling = ON
 Sampling rate = 35 samples per minute b. Click on the “Configure Spectrometer” button
 c. Click “Individual Wavelengths” from the drop-down menu and select the wavelength(s) that you would like to observe for your experiment. You do not have to observe the peak wavelength, but you should observe a wavelength that is near the peak.
3. Measuring the Reaction Time for Run #1 a. Remove Tubes 1A and 1B from the room temperature water bath and prepare to mix them. b. Click the “Collect” button to start the data collection at the same moment that you pour the contents of Tube 1A into Tube 1B. Pour the mixture back and forth a few times to mix thoroughly. Fill a cuvette with the mixture and place it in the SpectroVis Pro. c. Collect data until the absorbance levels out at or near zero. From the “Analyze” menu, choose
“Examine” to find the time that the reaction is complete. You do not need to print these graphs, just record the time at which the absorbance levels out near zero. If the color of your sample does not fade significantly in 2 minutes, inform your instructor or Lab Assistant.
4. Repeat your data collection of Run #1 until your times agree within 10% of each other.
Q: What is the difference between a reaction time and a reaction rate? You have measured reaction times, what do you have to do to determine the reaction rate for each run? Does a long reaction time correspond to a “fast” or a “slow” reaction rate?
D. Data Collection: Cold Run.
Once you have completed part C above and obtained reproducible reaction times you are ready to set up the cold run. Prepare samples “CA” and “CB” with the same reagents and volumes as “1A” and “1B” above. Prepare a separate blank (or reference tube) for the low temperature run, by adding about 8 mL of distilled water to a clean large test tube. Allow tubes CA, CB, and the blank to equilibrate in an ice/water bath (~0 ºC) for at least 5 minutes. Once the tubes are cold, quickly mix CA and CB. Record the time at the point of mixing to the nearest second and return the tube to the ice/water bath . Follow the course of the reaction by visually comparing the reaction solution with the blank. The Cold Run
NEVER GOES IN THE SPECTROVIS PRO ; you are monitoring the loss of color BY EYE . As the ice in the bath melts, continue to add ice and pour off water to keep the temperature at the bottom of the beaker constant. Record the time when the color due to I 2 has disappeared. The cold run takes a long time. Leave the system to react and continue on with the room temperature runs. Keep an eye on this cold run reaction mixture and record the time required for the reaction to reach completion.
Q: Why does the cold run take longer than the room temperature runs? Explain your answer on a molecular level.
13

E. Data Collection: Additional room temperature runs
You will be determining the order of this reaction with respect to the various reactants by comparing additional reactions to the “standard” reaction you performed above (Run #1). Design a series of additional experiments that you can use to determine the order of this reaction with respect to acetone, acid and iodine concentrations using the following guidelines:
1. Change only 1 variable at a time. Changing multiple variables will make it more difficult (although not necessarily impossible) to interpret your results.
2. The total volume of your sample should always be 8.00 mL.
3. Distilled/deionized water is present only to make the total volume of the solution 8.00mL. For example, if you choose to double the amount of acetone in the experiment to 4.00mL, then you should omit the water. (4.00 mL acetone(aq) + 2.00 mL HCl(aq) + 2.00 mL I 2 (aq) = 8.00 mL)
4. Consider the equipment that is already available in the lab. All measurements in your plan will have to be performed using the available equipment.
For each set of conditions, make repeat measurements until you are confident in your results. You will need at least 3 different sets of conditions to fully interpret your experiment. Your plan must be written down and checked by your instructor or Lab Assistant before you continue. Complete all of your room temperature reactions during the first week.
After completing all of your room temperature experiments, you may stop lab work for the first week. Before Week 2, calculate the concentrations of all reactants, calculate the reaction rates, and plan your experiments for next week.
F. Prepare a second “Cold” run using the same conditions as in week 1. While this is reacting, prepare a
“Cool” water bath at approximately 10 ºC. You may choose to monitor this bath temperature by using the temperature probe and a temperature vs. time data collection in LoggerPro. Stirring the bath with a magnetic stir bar, try to keep the temperature within ±0.5 ºC by adding an ice chip or two when the bath starts to warm up. Prepare at least 2 “Cool” runs using the same volumes as the “Cold” runs. Monitor these reactions for completion visually as you did with the “Cold” run.
III. Data Analysis: This experiment generates a LOT of data; it is probably best to organize this data in a table or multiple tables to keep it straight. There are also quite a few repetitive calculations in this experiment; you may wish to use MSExcel or another spreadsheet to perform your calculations.
A. Calculate the actual initial concentration of all reagents used in each reaction mixture.
Q: What is the difference between an “initial concentration” and a “stock concentration”? What is the limiting reagent in each of the experiments you performed? Is this consistent with your observations? What would you observe if the limiting reagent were different?
B. Calculate the average rate of each experiment from your reliable trials.
C. Qualitative determination of order with respect to each reactant.
Comparing your runs to Run #1, determine the order with respect to each reactant, assuming that the orders should (theoretically) be integers 0, 1, or 2. For example, if you double the concentration of iodine and the reaction rate is essentially unchanged, the order of the reaction with respect to iodine must be zero.
D. Quantitative determination of orders with respect to each reactant.
The orders with respect to each reactant can also be determined mathematically. In theory, the values should be integers, so determining their exact mathematic value is not always necessary, but it does allow a more complete analysis of the experimental error. Set up a ratio of the rate equation for two runs, cancel everything you can cancel, and use a logarithm to solve for the exponent (reaction order) that’s left over.
Q: Why is there error in your result? Create a list of sources of error in this experiment and estimate the extent of each source. For example, if the HCl(aq) stock solution is listed as “1.00M”, there is
14

error in that last digit; the “real” concentration is somewhere between 0.99M and 1.01M. That’s a
1% error. Is 1% a significant error in this experiment?
E. Calculate k, the rate law constant, for each of the runs, and the average value of k for all of your room temperature runs.
Q: How much error is there in your average room temperature k? How does this error impact the number of significant figures you can legitimately report in your value of k?
Q: How does your average room temperature value of k compare to the value of k for your cold run?
Is this difference consistent with the rates you observed for the reaction?
F. The activation energy, E a , may be determined using either the comparative form (for any 2 points) or the linear form (for all 3 points) of the Arrhenius equation. Calculate E a mathematically for each pair of temperatures (Room/Cool, Room/Cold, Cool/Cold), and graphically using all three temperatures. See your text for details.
Q: How do your pair-wise activation energies compare to one another? How do they compare to the graphically determined value? Which value do you prefer?
Q: Is this a “large” activation energy? How does the activation energy of a reaction influence the rate of the reaction? Did this reaction seem very fast, very slow or somewhere in the middle of the road?
15

Iron(III) Nitrate and Potassium Thiocyanate
Introduction:
Although equilibrium concentrations are very important quantities, they are often difficult to measure. If one of the products we are trying to examine is colored, it becomes significantly easier to determine its concentration in solution by measuring the absorbance of the solution and using Beer’s
Law.
Safety Concerns:
No specific concerns for this experiment.
Experimental Procedure:
I. Reaction between Iron (III) Nitrate and Potassium Thiocyanate
A. Dissolve a small amount (spatula tip) of solid iron (III) nitrate, Fe(NO 3 ) 3
·9H
2 O, in 20 mL of distilled water in a small beaker. Dissolve approximately the same amount of potassium thiocyanate solid, KSCN, in another 20 mL of water. Label each beaker and record the appearance of the solutions.
B. Mix about 10 mL of each of the prepared solutions in a different beaker. Record your observations. Save the remaining portions of the two solutions.
C. Assuming that this is an example of a metathesis reaction involving dissolved ions, which ions are responsible for the colored product you have observed? There is an assortment of different salts containing these ions in the lab. (Please leave the bottles where others can find them.)
Using these salts and the saved solutions from the preceding step, design and conduct several experiments that will enable you to conclude which ions are responsible for the formation of the colored complex. Record the results of your experiments. You must have a chemical test that shows whether each of the ions in the experiment {Fe
3+
(aq), NO 3
–
(aq), K
+
(aq), SCN
–
(aq)} is responsible for the observed reaction. Write the net-ionic equation for the reaction assuming a
1:1 stoichiometry for the reactants.
II Determination of the Beer’s Law Constant
A. Using pipettes, prepare a number of solutions containing 5.00mL of KSCN stock solution and
15.00 mL of Fe(NO 3 ) 3 (aq) stock solution. The results of the rest of this experiment will depend upon the Beer’s Law constant you calculate from these solutions, so prepare as many as necessary to give you confidence in your result. Allow these solutions to react for 5-10 minutes and then record their absorbance vs. wavelength spectra. Choose a wavelength at or near a peak in the spectrum and record in your lab notebook both the wavelength and the absorbance at this wavelength for the solutions you measured.
Q: What is the limiting reagent in these reactions? How will these conditions affect the equilibrium compared to a reaction that had no limiting reactant?
B. Assuming the limiting reactant is completely consumed to form the desired 1:1 product in these samples, calculate the value of “ε•l” in Beer’s Law at the wavelength you chose. Include error in your results.
Q: Is it a valid assumption that the limiting reactant is completely consumed in this reaction?
If so, what does this tell you about the magnitude of the equilibrium constant for this reaction? As an example, would this assumption be valid if the equilibrium constant was very small? Very large? Explain.
16

III. Exploring the equilibrium quantitatively
A. Obtain two burettes and label them as water and Fe(NO 3 ) 3 . Make sure the burettes are clean and they do not leak. In clean labeled beakers, take about 50 mL each of the stock KSCN and
Fe(NO 3 ) 3 solutions. Rinse each burette with several 3-5 mL portions of the solution that will be used in them subsequently. Fill the burettes with the respective solutions. Make sure that the stopcock and the burette tip do not have any air bubbles in them. If you have any questions, ask for help before you continue to prepare your samples.
B. You will be preparing at least 8 solutions (you may prepare more if you like), each with a total volume of approximately 20 mL. Prepare each sample in a clean dry large test tube. Each solution will contain 5.00 mL of KSCN solution {pipette}, 0.50-8.00 mL of Fe(NO 3 ) 3 (aq)
{burette}, and enough water {burette} to make approximately 20 mL of solution. For example, one of your samples might contain 5.00 mL KSCN(aq), 6.13 mL Fe(NO 3 ) 3 (aq), and 8.94 mL of water, for a total solution volume of 20.07 mL. It is not necessary for the total volume to be exactly 20.00 mL, but it is important that you know the total (additive) volume of the solution.
Your 8 samples should cover the full range of volumes of the iron solution listed above. If you report all of your solutions with a total volume of 20.00 mL, you are not using the equipment correctly and will receive zero points for this experiment.
C. Mix your samples thoroughly and allow them to react for 5-10 minutes before recording the spectra of each solution at the wavelength of maximum absorbance determined in part IIA. It is easiest to mix these solutions by covering the tube (rubber stopper or Parafilm) and inverting the tube several times .
Q: What is the limiting reagent in each of these samples? If you make the same assumption as in Part IIB above, what would you expect the absorbance of each solution to be? Is this what you observe? What does this observation tell you about the magnitude of the equilibrium constant for this reaction? It may be helpful to look at these data graphically.
Prepare a plot of absorbance vs. initial iron(III) concentration. Explain any trend you see in your data.
D. Calculate the initial concentrations {after mixing but before reacting} of the ions responsible for the observed reaction for each of the samples you have prepared. Using the value of “ε•l” calculated above, calculate the equilibrium concentration of the colored product in each of your samples. Once you know the initial concentration of each reactant and the equilibrium concentration of product, you should be able to calculate the equilibrium concentrations of all species and calculate the equilibrium constant for the reaction.
Q: Is the value of the equilibrium constant consistent with your predictions earlier in the experiment? Explain any deviations.
17

Introduction:
Acetic acid is one of the simplest and more common organic acids. It is the chemical responsible for the sour taste of vinegar. It is a weak acid and as such does not completely dissociate in water so the intact and dissociated forms are in equilibrium in solution. Because of this equilibrium, the percent ionization of acetic acid can be readily affected by subtle changes in experimental conditions. The percent ionization can be determined by measuring the pH of the test solutions to determine the concentration of H 3 O
+
(aq).
CH 3 CO 2 H(aq) + H 2 O(l) CH 3 CO 2
–
(aq) + H 3 O
+
(aq)
Safety Concerns:
Although the concentration of the acid used in this experiment is low and acetic acid is not exceptionally hazardous, caution is still required. Wear your goggles at all times in the lab. Wipe up any spills immediately and rinse the spilled area with water.
Experimental Procedure:
Part 1 – The Effect of Water on pH and Ionization
1A. In the laboratory you will find a stock solution of acetic acid that has a formal concentration of about 1 M. Obtain 20-25 mL of this stock solution in a clean, dry beaker. Using a 10.00 mL pipette* and a 100.0 mL volumetric flask prepare a solution that has a concentration one-tenth that of the stock solution. Store this acetic acid solution in a small, clean, dry beaker and label it,
Solution #1. Thoroughly rinse (it is not necessary to dry) the volumetric flask with water for use in part B.
* Be sure that the pipette is clean and rinse it with several small volumes of the solution to be transferred so that the solution is not accidentally diluted because of any water in the pipette.
1B. Using solution #1, prepare a second solution (Solution #2) that is one-tenth as concentrated as solution #1 (one-hundredth as concentrated as the stock solution). Again be careful to properly rinse the pipette with solution #1 prior to pipetting. Calculate the concentration of this solution and label it Solution #2.
1C. Using solution #2 prepare two more solutions: one that is one-fifth as concentrated as solution #2 labeled Solution #3 (one-five hundredth as conc. as the stock), and one that is one-tenth as concentrated as solution #2 labeled Solution #4 (one-thousandth as conc. as the stock). Calculate the exact formal concentration of these solutions.
1D. Measure the pH of each solution. Save all solutions for now.
The first solution (one-tenth the concentration of the stock solution or about 0.1 M acetic acid) will be used in part II.
Q: Were the measured pH readings of these solutions different? Is there a trend? Can you explain the trend in the measured pH?
1E. Use the measured pH and the formal concentration of the acid to determine the actual concentrations of the species H 3 O
+
, CH 3 CO 2
–
and CH 3 CO 2 H. Calculate the % ionization of
CH
3
CO
2
H in each solution. You may wish to consult other groups regarding your values.
Q: Explain the trend in your calculated percent ionization values. Are the concentrations of
H 3 O
+
, CH 3 CO 2
–
and CH 3 CO 2 H related? Is there a mathematical expression that describes this relationship?
18

Part 2 – The Effect of Sodium Acetate on pH and Ionization
Sodium acetate, CH 3 COONa, is a soluble salt.
Q: What ions will this salt produce in an aqueous solution? Predict whether the addition of sodium acetate will have an impact on the reaction of acetic acid with water shown on the previous page? Do you expect the pH of the acetic acid solution to change if sodium acetate is added to the solution?
2A. Weigh out about 0.1 g of anhydrous sodium acetate, CH
3
COONa. Record the mass in your lab notebook. Using a graduated cylinder, measure 50.0 mL of the approximately 0.1 M acetic acid solution saved from Part 1 into a clean, dry 150 mL beaker. Place a Teflon-coated stirring bar in the beaker and use a stirring motor to stir gently. (Don’t make a tornado!) Place the pH probe in the solution (be careful to not hit the probe with the stir bar). When the pH reading has stabilized, record the pH value. Add the first sample of sodium acetate (solid). When the sample has completely dissolved and the pH reading has stabilized again, record the pH value.
Q: Did the pH reading change as you predicted it would when you added the sodium acetate to the acetic acid solution? Can you explain the change in the measured pH?
Based on the change in pH when sodium acetate was added how would the percent ionization value change? Do you think that the mathematical relationship between the concentrations of H
3
O
+
, CH
3
CO
2
–
and CH
3
CO
2
H that you obtained in part 1 will still apply in this solution with added sodium acetate? pH and the Concentration of Acids and Bases
The acidity (or basicity) of a solution can be determined by measuring the concentration of hydronium ions, H
3
O
+
(aq), that result from the acid donating H
+
(aq) to a water molecule. Because these concentrations are often quite small, it is more convenient to express them as pH. The pH of a solution is the negative logarithm of [H
3
O
+
]: pH = –log[H
[H
3
O
+
] = 10
3
O
+
–pH
]
It is important to remember that the measured pH of a solution indicates how acidic or basic the solution is, it does not directly indicate the strength of the acid (or base). The strength of an acid depends upon its ability to ionize, usually in the presence of water. This can be described by the equation:
HA(aq) + H
2 where “HA(aq)” is an acid and “A –
O(l) H
3
O
+
(aq) + A
–
(aq)
(aq)” is its conjugate base. The strength of the acid is determined by the position of the equilibrium; the more product-favored the equilibrium, the “stronger” the acid.
Acids are typically classified as “weak” if the equilibrium is reactant-favored and “strong” if the equilibrium is product-favored.
The strength of an acid can similarly be expressed as its percent ionization. Similar to the equilibrium constant, the higher an acid’s percent ionization, the “stronger” the acid. A “strong” acid is considered to be (essentially) 100% ionized in solution. Percent ionization compares the amount of acid that has ionized to the total amount of acid present. The “total amount of acid present” expressed as a concentration is often called the “formal concentration” of the acid and is usually the “initial concentration” in equilibrium problems.
19

In this experiment you will be given a solution of a base, Na 2 CO 3 (aq), and a solution of an acid,
HCl(aq), but the concentration of only one of these solutions will be known. The goal of this experiment is to determine the concentration of the unknown solution using an appropriate set of data obtained from titrations using three different acid-base (visual) indicator solutions and a pH meter. Visual indicators change color over a relatively narrow pH range, known as the endpoint.
When using a visual pH indicator it is important to match the endpoint of the indicator with the expected pH of the equivalence point of the titration being observed. The equivalence point is the point in a titration when the moles of acid and base present in the reaction match the stoichiometry of an appropriately balanced chemical equation.
Safety Concerns:
The solutions of acids and bases used in this experiment are dilute, however, caution is still required.
Wear goggles at all times in the lab. Wipe up any spills immediately and rinse the area with water.
Experimental Procedure:
IA. Obtain a burette and make sure that it is clean and does not leak. If necessary, clean the burette with a burette brush and soapy water and rinse it with distilled water. Then rinse it several times with a few milliliters of the stock HCl solution, make sure that the stopcock and burette tip are also thoroughly rinsed with HCl. Fill the burette with the HCl solution making sure that there are no air bubbles and no leaks in the stopcock or tip.
IB. Obtain about 100 mL of stock Na 2 CO 3 solution in a clean, dry beaker. Rinse a clean 20 mL pipette with a few milliliters of Na
2
CO
3
solution. Then pipette 20.00 mL of the Na
2
CO
3 solution into a clean, 125 mL Erlenmeyer flask. Add 2 drops of phenolphthalein indicator solution to the Na 2 CO 3 solution. Swirl to mix.
IC. Slowly allow the HCl solution to drain into the Na
2
CO
3
solution while swirling the flask. As the endpoint approaches, endpoint color persists longer and longer. For multi-color indicators, the correct endpoint color is the intermediate color between the two extremes. Slow down the drain rate as you approach the endpoint. As the endpoint gets closer add HCl drop by drop, swirling the reaction mixture well before the next addition. Stop the addition of HCl as soon as the indicator has changed color. Record the volume of HCl necessary to give an indicator color change that is permanent. If you feel that you overshot the endpoint repeat the experiment.
The color changes to expect when going from a basic to an acidic solution are:
Phenolphthalein pink

colorless
Bromothymol blue
Methyl orange blue

(green)

yellow yellow

(orange)

red
ID. Repeat the titration-using bromothymol blue as the indicator.
IE. Repeat the titration-using methyl-orange as the indicator.
Q: Did all three indicators change color when the same volume of acid was added? Are the volumes related? Did all of the indicators have a sharp, easy to detect endpoint?
II. pH Titration: Set up the apparatus as shown in the diagram.
Standardize the pH meter using pH 4 and 7 buffers. Consult the instructor for assistance, if needed. In this part of the pH sensing device experiment use a pH meter to monitor the pH variation during titration. Use a 150 mL beaker, as there is not adequate room to place the pH probe and the burette tip into of an Erlenmeyer flask. Pipette 20.0 mL of
20

Na 2 CO 3 (aq) stock solution into the beaker and add ~25 mL of distilled water to the pipetted sample. Also add 2 drops of methyl orange indicator to this reaction system. Use a magnetic stirring bar to mix the solution efficiently.
The rate of change of the pH with the addition of HCl will vary considerably during the titration. Run a titration by adding ~0.5 mL aliquots of HCl. Allow a few seconds of mixing time before you take readings. Record the pH and exact volume of added HCl during the titration. Also record the pH and volume of HCl added when the indicator changes its color.
Continue the titration until the pH reaches a value of 2, which will be well past the endpoint.
Copy the graph using the Snipping Tool accessory and paste it into a Microsoft Word document and then email this file to each member of your group. You will need to include and make reference to the pH titration in your hand-in.
Q: Explain the relationship between the two equivalence point volumes in a chemically reasonable way. Given the observed titration curve, review your data from the indicator titrations and explain your observations. Are all three visual indicators appropriate for this titration? Why or why not? How does your calculation of the unknown concentration change when using each visual indicator? Calculate the unknown concentration using as many endpoints/equivalence points as you feel are valid.
III. Buffers
Graduated cylinders may be used for measurements in this part for simplicity. When half an equivalent of acid has been added to a sample of carbonate ions, the concentration of the species CO 3
2–
and HCO 3
–
are equal which means that you have produced a good buffer.
IIIA. In a 150 mL beaker, prepare a buffer solution in which the concentration of CO 3
2–
and HCO 3
– are equal by starting with 50.0 mL of the stock carbonate solution and adding an appropriate volume of HCl(aq) to protonate half of the carbonate ions present in solution. This is the concentrated buffer solution. Transfer 25 mL of this concentrated buffer solution to a second beaker and add 40.0 mL of deionized water to make a dilute buffer solution, mix well. You will need two 15 mL samples of each buffer in separate beakers (4 total samples) for testing in the next part of the experiment.
IIIB. Measure the pH of one of the two concentrated buffer solutions. Add 1drop of 1M HCl and record the new pH of the solution. To the second concentrated buffer solution add 1drop of
1M NaOH and record the new pH of the solution.
IIIC. Repeat IIIB for the two dilute buffer solutions.
IIID. Repeat IIIB with two 15mL samples of distilled water with a spatula tip of NaCl added. {This is a blank for reference. The NaCl is necessary to allow conductivity in the water so the pH meter can operate properly but the NaCl does not affect the pH}
Q:
Were both buffers (concentrated and dilute) equally “good” at limiting the change in pH when strong acid and/or base are added? Explain. Over what range of pH would you expect this buffer system to be “good”? Explain.
21

IV. Indicator Spectra
Add 50.0 mL of deionized water, 5.0 mL of Na 2 CO 3 (aq) stock solution, and 10-15 drops of bromothymol blue indicator to a 150 mL beaker and stir gently with a magnetic stir bar.
OPEN UP A SEPARATE COPY OF LOGGERPRO and set up to record an Absorbance vs.
Wavelength spectrum. Fill a cuvette with your carbonate solution and record the Absorbance vs. Wavelength spectrum. Return this solution to the beaker , add 1-2 drops of HCl(aq) to the beaker, let the solution stir for a few seconds, and measure the Absorbance vs. Wavelength spectrum of the solution again by taking some of the solution out and putting it in a cuvette.
Keep all of the spectra on the screen for this entire portion of the experiment so you can see how the spectrum changes.
Return the cuvette solution to the beaker and add 1-2 drops of
HCl(aq), mix well, and measure the Absorbance vs. Wavelength spectrum. Continue this procedure until you have added 4-6 drops of HCl(aq) beyond the endpoint . Remove the colored background from the spectrum and copy using the Snipping Tool accessory. Paste the spectrum into a Microsoft Word document and email the file to each member of your group.
Q: How did the Absorbance vs. Wavelength spectrum change throughout this portion of the experiment?
22

3+
2+
2+
2+
+
How do you know what is in a sample? Modern instrumental methods allow the composition of a sample to be determined quite accurately by simply injecting some of the sample into an instrument, but many of these instruments are fairly slow and large which prevents their use outside a laboratory setting. How can a geologist determine the type of rock she finds in the field? How can a police officer decide whether the white powder found in a suspect’s car is baby powder or illegal drugs?
Before the advent of modern instrumentation, chemists were forced to rely upon their observations of physical properties and reactivity, and had to use those observations in a logical progression to determine the composition of an unknown sample. In this experiment, we will be exploring the behavior of a variety of metal cations (positively charged ions). Using our observations, we will identify an unknown sample
You have been provided with solutions of the nitrate salts of the 5 cations listed in the title of the experiment. You will test each of these cations under a variety of conditions as outlined below to observe their behavior under a variety of conditions. Your observations should include:
1. Formation of a precipitate
2. Color of any precipitate that may form
3. Evolution of a gas
4. Dissolution of a precipitate
5. Changes in the color of the solution
Your ability to correctly identify your unknown is dependent upon your observations, so record everything you see! Your observations will also depend very strongly upon the quality of the reagents you use to perform these tests. The supplied reagents are clean, but it is very easy to crosscontaminate them with poor lab procedure. When dispensing reagents into a test tube with a dropper, the tip of the dropper must NEVER touch the test tube. Always release the drop from a little above the mouth of the test tube.
Safety Concerns: Most of the chemicals we are using today are not hazardous; however, you must always exercise caution when performing any experiment.
You must ALWAYS wear your safety goggles while the lab. Even if you aren’t performing an experiment, your neighbor might be and there is always the possibility that a chemical can splash toward your face. ALWAYS wear your goggles!
Strong acids and bases can burn your skin, but all of the acids and bases we will be using can be safely washed off with water. If you get any reagent on your skin, wash the area thoroughly with water at the sink.
Both ammonia and hydrochloric acid have VERY strong odors. When using these reagents, avoid smelling them as much as possible
Experimental Procedure:
You do not have to do these tests in numerical order, but once you start doing a number, do all parts of that number in order. For example, you can do #5 first, but you must do 5a then 5b then 5c.
1. To 10 drops of fresh test solutions, add 10 drops of 15M NH
3
(aq) dropwise.
2. To 5 drops of fresh test solutions, add 2-3 drops of K
2
CrO
4
(aq).
3a. Put 10 drops of each cation solution into small test tubes and add 5 drops of 3M H
2
SO
4
to each.
Record any observations.
23

3b. If a precipitate forms, centrifuge the mixture. Decant the supernatant; that is, carefully pour the supernatant out of the tube leaving the precipitate behind. Wash the precipitate once by adding about 2mL of distilled water, mixing, centrifuging and discarding the wash liquid. To the precipitate, add 6M HNO
3
(10-20 drops) to see if the precipitate dissolves.
4a. To 10 drops of fresh test solutions, add 5 drops of 6M HCl. If a precipitate forms, mix the sample thoroughly, centrifuge, and discard the supernatant. Wash the precipitate with about 2mL of water, centrifuge and discard the supernatant.
4b. To any precipitates formed in 4a, add about 3mL of fresh distilled water. Heat the tube in a hot sand bath until it boils gently. If a precipitate remains when the sample is hot, quickly centrifuge and decant the hot liquid into a separate tube. (Save the precipitate for Part 4d.)
4c. To the hot liquid that was decanted in Part 4b or to any solutions from Part 4b where the solid completely dissolved, add 1-2 drops of K
2
CrO
4
(aq) and mix thoroughly.
4d. Wash any precipitate remaining from Part 4b with about 2mL of distilled water. Centrifuge and discard the wash liquid. Add 5-10 drops of 15M NH
3
(aq) to the precipitate, mix well and observe any changes in the precipitate and the supernatant.
5a. To 10 drops of fresh test solutions, add 10 drops of 3M NaOH one drop at a time, mixing and observing the sample carefully after every drop . Centrifuge any samples that have precipitates after all 10 drops are added, discard the supernatant and save the precipitate for 5b.
5b. Wash the precipitates from Part 5a once with water, centrifuge and discard the supernatant. To the precipitate, add 10 drops of 15M NH
3
(aq) and mix thoroughly.
6a. To 10 drops of fresh test solutions, add 10 drops of 3M NaOH. Centrifuge and decant the supernatant, saving the precipitates for Part 6b.
6b. Add 5-10 drops of 6M HNO
3
to the precipitate and mix thoroughly.
6c. Add 2 drops of 1M NH
4
SCN solution to the solutions from Part 6b.
Analysis of an Unknown:
The unknown you are given will contain a mixture of the metal cations you have tested. Given the results and observations you have collected with the known samples, devise a plan (a flow chart), which will allow you to separate and identify each of the metal cations that are present in your mixture. Your unknown could contain as few as 1 cation or as many as 5, so your flow chart should allow you to independently and positively verify the identity of each of the 5 cations and physically separate these 5 cations from each other . This means that in many cases negative results will not be useful, so be very careful! Your flow chart should be recorded in your lab notebook.
Chemical Equations:
Give net-ionic equations to represent only those species that are reacting in each case. Be sure that your equations are balanced. The following guidelines will enable you to make accurate predictions of the products in most cases.
1. Assume that most reactions result from a simple combination of cations and anions.
2. If the product is a solid assume that it is electrically neutral, that is, the total cationic(positive) and anionic(negative) charges in the solid compound are equal.
3. If you have made an observation that indicates that a reaction has occurred and the product is in solution assume that it is not necessarily electrically neutral but may possess a net charge, which may be either positive or negative.
24

Introduction:
The vast majority of chemistry involves the movement and redistribution of electrons. In many systems, this happens without any consequence, but if the system can be constructed such that the moving electrons are passed through an external circuit, that electron movement can be harnessed to do useful work. This is the basis of all batteries and other power sources. In this experiment, we will be exploring the behavior of various electrochemical half-cells and using our observations to rank their activity.
Safety Concerns:
The solutions used in this experiment should be handled with caution. Wear your goggles at all times in the lab. Rinse your hands thoroughly if any solution splashes or spills on you. Wipe up any spills immediately and rinse the spilled area with water. Do not pour any solutions down the drain, but instead, place them in the appropriately labeled waste containers in the hood.
Experimental Procedure:
I. Reactions of metals with metal ions
Obtain a strip of Zn, Cu, Pb, and Ag metals. If necessary, clean the surface of pieces of metal using sandpaper or steel wool. { Do not clean electrodes directly on the benchtop!
} On each metal strip, place a single drop of each metal ion solution (don’t let them touch!) and determine whether or not a reaction has happened. It may be useful to put a drop of water on each strip to serve as a reference.
Q: Which metal is most reactive? Least reactive? Describe how you determine their reactivity and rank all 4 metals from most reactive to least reactive.
Q: Which metal ion is most reactive? Least reactive? Describe how you determine their reactivity and rank all 4 metal ions from most reactive to least reactive.
Q: Are your rankings related? Offer an explanation of any relationship you observe.
II. Voltaic Cells
IIA. Fill four clean small test tubes half full of the solutions of the same four metal ions studied in part I. Fill another small test tube half full of a Ni(NO
3
)
2
solution. Label the tubes appropriately. Place a clean strip of each metal in the solution of its corresponding ion. Each of these tubes is a M +x |M-type half-cell. Coupling two half-cells with a salt bridge forms an electrochemical cell with a multimeter completing the external circuit.
IIB. Measure the cell potential of each electrochemical cell shown in the table. Connect the multimeter leads firmly on the strips of metal.
Do not simply touch the leads to the metal electrodes but connect them firmly to the
Cell Cell
Zn | Zn 2+ || Cu 2+ | Cu Pb | Pb 2+ || Ni 2+ | Ni
Zn | Zn 2+ || Ni 2+ | Ni Pb | Pb 2+ || Ag + | Ag metal. It often helps to clean the surface of
Zn | Zn 2+ || Ag + | Ag Ag | Ag + || Cu 2+ | Cu the electrode first for better electrical contact.
The current passing through these cells is so low that there is no notable danger involved.
Zn | Zn 2+ || Pb 2+ | Pb Ag | Ag + || Ni 2+ | Ni
Pb | Pb 2+ || Cu 2+ | Cu Ni | Ni 2+ || Cu 2+ | Cu
It is a good practice to connect the voltmeter so that the multimeter negative (black lead) is connected to the left-hand side electrode of the cells shown in the table. Being consistent will make it MUCH easier to interpret your results.
Q: What does it mean if the measured potential is negative? What does this tell you about the cells as written in the table?
25

IIC. Pick one of the electrochemical cells from the previous section and reassemble it with the multimeter hooked up backwards. That is, connect the negative black lead to the right hand electrode.
Q: How does the new measurement compare to the potential you measured in Part IIB?
Explain any difference.
IID. Starting with your most reactive metal as determined in Part I, note the measured potential between it and the other three metals used in Part I. You can re-measure the potentials if you are unsure of your interpretation of the results.
Q: Is there a relationship between the ranking you determined in Part I and the potentials you measured for each pair of metals?
IIE. Compare the measured cell potentials between various pairs of metals, both those that are adjacent to each other in the ranking from Part I and those that are not. Again, you can remeasure the potentials if you are unsure of your interpretation of the results. For example, if
M1 is more reactive than M2, and M2 is more reactive than M3, look at your measured cell potentials for M1-M2, M2-M3 and M1-M3.
Q: Is there a relationship between the measured cell potentials for adjacent metals and the measured cell potentials for non-adjacent pairs?
Q: Using all of your measured cell potentials, where would you place nickel in your ranking?
Explain how your data supports this ranking.
III. Concentration Effects
IIIA. In this part of the experiment the Cu +2 (aq)|Cu half-cell will be coupled with several Ag +
(aq)|Ag half-cells in which the concentration of Ag + (aq) is varied systematically. Dilute 1.0 mL of the 0.1 M Ag + (aq) solution to 10.0 mL with distilled water to prepare a 0.01 M Ag + (aq) solution. You can use a graduated cylinder for this dilution; mix the dilution well and transfer to a large test tube. Perform two more serial dilutions in the same manner in order to prepare
0.001 M Ag + (aq) and 0.0001 M Ag + (aq) solutions. Place a few milliliters of each Ag + (aq) solution in separate clean, small test tubes.
IIIB. In the tube containing the most dilute solution {0.0001 M Ag + (aq)} place a clean strip of Ag.
Couple this half-cell to the Cu +2 | Cu half-cell and measure the cell potential of the spontaneous reaction.
IIIC. Remove the salt bridge and Ag metal strip from the Ag + (aq) solution, dab the salt bridge on a paper towel and then place the salt bridge and Ag metal strip into the 0.0010 M Ag + (aq) solution. Couple this half-cell with the copper half-cell, and measure and record the cell voltage.
IIID. Measure the cell potential with the 0.010 M Ag + (aq), and 0.10 M Ag + (aq) solutions in the same manner.
IIIE. How does the cell potential change when the concentration of Ag + (aq) in the half-cells changes? Prepare graphs of E cell
vs. [Ag + ] and E cell
vs. log[Ag + ]. Use E cell
for the spontaneous reaction in your graphs.
Q: How would you describe the relationship (shape of the graph) between cell potential and concentration?
Q: How would you describe the relationship (shape of the graph) between cell potential and log(concentration)?
Q: Which of these graphs shows a clearer relationship between cell voltages and concentration of Ag + (aq) ions?
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