Lab Notebook and Grading Scheme for Analytical Chemistry II - Spring 2018, Summaries of Analytical Chemistry

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CHEM 232 ANALYTICAL CHEMISTRY LAB II

Laboratory Manual

Department of Chemistry

İZMİR INSTITUTE OF TECHNOLOGY

Urla / İZMİR

Urla – İZMİR

Spring 2018

CHEM 232 ANALYTICAL CHEMISTRY LABORATORY II

This manual has been prepared for the CHEM 232 Analytical Chemistry Laboratory

and includes the experiments, which are related to the topics covered in the CHEM 202

Analytical Chemistry II course. The main purpose of this laboratory is to provide the

students an appreciation for the potential applications and limitations of analytical methods

of analysis. It is also aimed to provide the students an opportunity to develop their abilities

in the laboratory skills required for accurate and precise chemical analyses. Therefore, it is

expected that with an acceptable degree of proficiency the students will grasp and apply the

principles of analytical chemistry to obtain reliable results in the laboratory experiments.

COURSE POLICIES AND INFORMATION

In this laboratory, you will be working as a team with three or four persons in each

group. Each project will take two or three weeks depending on the project given to you.

During the first lab period your instructor will assign you to a group. Each student in the

group must have a Laboratory Notebook and bring it to the laboratory every week. You

should keep a good notebook with all the calculations and results, because your instructors

will grade your Lab Notebooks at the end of each experiment. You are also expected to

write a purpose and procedure for each experiment prior to coming to the laboratory. At the

end of each lab period you will have two days to complete your notebook. THE

NOTEBOOK MUST BE COMPLETED AND TURNED INTO THE TEACHING

ASSISTANT WITHIN TWO DAYS AFTER COMPLETION OF THE LAB PERIOD.

Also there will be a pre-lab quiz given at the beginning of each lab period. Therefore, you

must be prepared for the lab experiment before it begins.

Your overall lab grade will be calculated by taking the sum of all your individual

project grades and dividing it by the total number of lab experiments. There will also be a

final exam at the end of the semester.

Attendance is required and you are expected to attend all scheduled laboratory

sessions. If you miss more than two-lab sessions without a valid reason you will be

automatically dropped from lab. Should you be sick and can not come to lab, you need to

bring a doctor’s excuse from the UNIVERSITY HEALTH CENTER or an official

3. Towards the end of the lab experiment, the following must be in your lab notebook (4 points):

H. Problems encountered, proposed solutions for problems.

I. Sources of error, estimates or precision.

J. Another student in your group or your TA must look briefly at your notes while you explain

what you did, he or she must print their name, sign their name and also write the date at the

bottom of each page used that day.

4. After the experiment is finished, include the following (6 points):

K. An analysis of the results (calculations),

L. A brief interpretation of the results,

M. Ideas for future experimentation to complement this experiment, references.

5. Also the lab notebook will be graded for legibility (2 pts).

The names of all students who also directly helped must be included in the lab notebook.

Only write the names of those persons from whom you received help.

Sometimes data comes from instrumentation in the form of printouts or electronic copies.

These must be handled correctly with the notebook. Original raw results such as spectra, etc. must be

stapled in the appropriate location in the notebook. For those experiments where only one copy of

results is available, each student must obtain a photocopy of this raw result and staple it in the

appropriate location in his or her own notebook. These spectra must be signed, witnessed, and

dated also.

Finally, at the discretion of the TA or the teacher, notebooks may be graded according to

additional criteria specific to each experiment.

Note the following format for citing the lab manual or handout [5], an internet

reference [1] a book without an editor [2], a book with an editor [3], and a journal article [4]

at the end of your reports.

1. Smith College Lab Report

http://www.science.smith.edu/departments/Chem/Courses/labreports.htm

( accessed August 21, 2001 ).

2. Skoog, D.A.; West, D.M.; Holler, F.J.; Crouch, S.R. Fundamentals of Analytical

Chemistry 8 th ed.;

Thomson Brooks/Cole: Belmont, CA, 2004 ; pp 314-336.

3. Almlof, J.; Gropen, O. Relativistic “Effects in Chemistry.” In Reviews in Computational

Chemistry;

Lipkowitz, K.B.; Boyd, D.B., Eds. VCH: New York, 1996 ; Vol. 8 pp. 206-210.

4. Klinger, J. “Influence of Pretreatment on Sodium Powder.” Chem Mater. 2005 , 17, 2755-

6. Chemistry 215 Analytical Chemistry Laboratory Manual. 2009. “Gravimetric

Determination of Chloride and Sulfate in Some Inorganic Salts,” pp 11-16. Izmir

Institute of Technology, Turkey.

FORMAT FOR ORAL PRESENTATIONS

You will be required to give one oral presentation during this semester. In addition

to any guidelines given by the teaching assistants you must be aware of the following:

1. Any presentation that is suspected of having been copied from another source

(another student, the internet, etc.) will be investigated and graded accordingly. No copying

is allowed. A copied presentation will receive a grade of zero points.

2. For EACH slide, any information that is not general information must be properly

referenced AT THE BOTTOM OF THAT SLIDE. The referenced source must be written in

the proper format and consistent throughout the presentation.

3. The presentation must represent your own attempt to summarize the work of

others and you must place references properly for each slide.

4. After presenting the information, you are encouraged to include any ideas of your

own. Creativity and originality are encouraged.

5. Slides must be prepared using PowerPoint and your rough draft must be checked

prior to presentation by your teaching assistance.

6. At least the following slides must be included:

a. Title (with your name, address (IYTE Chemistry Depart.), and date)

b. Background and Importance of the topic

c. Previous research and findings on/about the topic

d. Current research and findings on/about the topic

e. Future goals of research related to the topic

f. Personal ideas or directions you may have for the research

g. Conclusions or Summary

h. Acknowledgements

The presentation should not be over 10 minutes. Two minutes will be allowed for questions.

GRADING SCHEME FOR THE COURSE:

Oral presentation 10% (individual grade)

Lab Quizzes 25 % (individual grade)

Lab. Notebook 25 % (individual grade)

Lab. Final Exam 30 % (individual grade)

Participation 10% (individual grade)

bit of the vapor toward your nose. Do not stick your nose in and inhale vapor directly from

the test tube. Always wash your hands before leaving the laboratory.

Eating and drinking any type of food are prohibited in the laboratory at all times.

Smoking is not allowed. Anyone who refuses to do so will be forced to leave the laboratory.

Clothing and Footwear : Everyone must wear a lab coat during the lab and no shorts

and sandals are allowed. Students who come to lab without proper clotting and shoes will be

asked to go back to change his or her clothing. If they do not come on time they will be

counted as an absent. Long hair should be securely tied back to avoid the risk of being set on

fire. If large amounts of chemicals are spilled on your body, immediately remove the

contaminated clothing and use the safety shower if available. Make sure to inform your

instructor about the problem. Do not leave your coats and back packs on the benchs because

they may be contaminated. No headphones, Walkmans, mp3 players or cell phones are

allowed in the lab because they interfere with your ability to hear what is going on in the

Lab. Cell phones must be turned off.

Fire: In case of fire or an accident, inform your instructor at once. Note the location

of fire extinguishers and, if available, safety showers and safety blankets as soon as you

enter the laboratory so that you may use them if needed. Never perform an unauthorized

experiment in the laboratory. Never assume that it is not necessary to inform your instructor

for small accidents. Notify him/her no matter how slight it is.

Laboratory Care and Waste Disposal

Remember that the equipment you use in this laboratory will also be used by many

other students. Please leave the equipment and all workspaces as you wish to find them.

After the end of the each lab, clean off your work area. Wash your glassware. When

weighing any material on the balances, do not weigh directly onto the balance pan. Weigh

your material on a piece of weighing paper. The balances are very sensitive instruments and

should be treated with great care.

If you take more reagents than you need, do not put excess back into the bottle. It

may be contaminated. Threat it as waste and dispose of it accordingly. It is most likely that,

during any experiment you will perform, you will generate some waste chemicals and

solutions to dispose of. Never put them down the sink unless specifically told to do so by

your instructor. There will be inorganic, organic, and solid waste containers in the lab.

Dispose of your waste in the appropriate container.

Analytical Chemistry II - TIME SCHEDULE OF THE EXPERIMENTS - Spring 2018

Teaching Assistants: Emre Yusuf GÖL, Seçil SEVİM ÜNLÜTÜRK, Seray Ece

KESKİN, Ahmet AYTEKİN

March 7 Introduction, safety, items required to enter lab, expectations

March 14 Lab 1 : Determination of Ascorbic Acid in Orange Juice

Drawing (raffle) will be held for your general topic for oral presentation (volumetry, gravimetry or

potentiometry).

March 21 Lab 2 : Standardization of Permanganate Solution

March 28 Lab 3: Determination of Calcium in Limestone

April 4 Lab 4: Determination of Hardness in Water Samples with EDTA

******Turn in a summary of your proposal for oral presentation.**

April 11 Lab 5: Potentiometric Titration of Phosphoric Acid

April 18 Lab 6: Determination of Equilibrium Constants for Complex Ions of Silver

April 25 Lab 7: Potentiometric Titration of Chloride and Iodide Mixtures

May 2 Lab 8: The Potentiometric Determination of Solute Species in a Carbonate Mixture

May 9 ******* Oral Presentations *******

May 16 ******* Oral Presentations *******

May 23 Makeup (last possible date - to be used at the discretion of the TAs)

Final Lab Exam will be officially announced and given during the exam week.

Experiment 2. Standardization of Permanganate Solution

Preparation of 0.02 M Potassium Permanganate:

PROCEDURE:

Dissolve about 3.2 g of KMnO 4 in 1 L of distilled water. Keep the solution at a gentle boil for about 1 hr.

Cover and let stand overnight. Remove MnO 2 by filtration (Note 1) through a fine-porosity filtering crucible

(Note 2) or through a Gooch crucible fitted with glass mats. Transfer the solution to a clean glass stoppered

bottle; store in the dark when not in use.

Notes

1. Heating and filtering can be omitted if the permanganate solution is standardized and used on the same

day.

2. Remove the MnO 2 that collects on the fritted plate with 1^ M H 2 SO 4 containing a few milliliters of 3%

H 2 O 2 , followed by a rinse with copious quantities of water

Standardization of Potassium Permanganate Solutions

PROCEDURE:

Dry about 1.5 g of primary-standard Na 2 C 2 O 4 at 110°C for at least 1 hr. Cool in a desiccator;

weigh (to the nearest 0.1 mg) individual 0.2-g to 0.3-g samples into 400-mL beakers. Dissolve

each in about 250 mL of 1 M H 2 SO 4. Heat each solution to 80°C to 90°C, and titrate with KMnO 4

while stirring with a thermometer. The pink color imparted by one addition should be permitted to

disappear before any further titrant is introduced (Notes 1 and 2). Reheat if the temperature drops

below 60°C. Take the first persistent (30 s) pink color as the end point (Notes 3 and 4). Determine

a blank by titrating an equal volume of the 1 M H 2 SO 4. Correct the titration data for the blank, and

calculate the concentration of the permanganate solution (Note 5).

Notes

1. Promptly wash any KMnO 4 that spatters on the walls of the beaker into the bulk of the

liquid with a stream of water.

2. Finely divided MnO 2 will form along with Mn

2+ if the KMnO 4 is added too rapidly, and it

will cause the solution to acquire a faint brown discoloration. Precipitate formation is not a serious

problem so long as sufficient oxalate remains to reduce the MnO 2 to Mn

2+ ; the titration is simply

discontinued until the brown color disappears. The solution must be free of MnO 2 at the end point.

3. (^) The surface of the permanganate solution rather than the bottom of the meniscus can be

used to measure titrant volumes. Alternatively, backlighting with a flashlight or a match will

permit reading of the meniscus in the conventional manner.

4. (^) A permanganate solution should not be allowed to stand in a buret any longer than

necessary because partial decomposition to MnO 2 may occur. Freshly formed MnO 2 can be

removed from a glass surface with 1 M H 2 SO 4 containing a small amount of 3% H 2 O 2.

5. (^) This procedure yields molarities that are a few tenths of a percent low. For more accurate

results, introduce from a buret sufficient permanganate to react with 90% to 95% of the oxalate

(about 40 mL of 0.02 M KMnO 4 for a 0.3-g sample). Let the solution stand until the permanganate

color disappears. Then warm to about 60°C and complete the titration, taking the first permanent

pink (30 s) as the end point. Determine a blank by titrating an equal volume of the 1 M H 2 SO 4.

method that follows is remarkably effective for determining calcium in most limestones. Iron and

aluminum, in amounts equivalent to that of calcium, do not interfere. Small amounts of manganese

and titanium can also be tolerated.

PROCEDURE:

Sample Preparation

Dry the unknown for 1 to 2 hr at 110°C, and cool in a desiccator. If the material is readily

decomposed in acid, weigh 3 portions of the sample each 0.25-g to 0.30-g samples (to the nearest 0.1 mg)

into 250-mL beakers. Add 10 mL of water to each sample and cover with a watch glass. Add 10 mL of

concentrated HCl dropwise, taking care to avoid losses due to spattering as the acid is introduced.

Precipitation of Calcium Oxalate

Dilute each sample solution to about 50 mL, heat to boiling, and add 100 mL of hot 6% (w/v)

(NH 4 ) 2 C 2 O 4 solution. Add 3 to 4 drops of methyl red, and precipitate CaC 2 O 4 by slowly adding 6 M NH 3.

As the indicator starts to change color, add the NH 3 at a rate of one drop every 3 to 4 s. Continue until the

solutions become the intermediate yellow-orange color of the indicator (pH 4.5 to 5.5). Allow the solutions

to stand for no more than 30 min (Note) and filter; medium-porosity filtering crucibles or Gooch crucibles

with glass mats are satisfactory. Wash the precipitates with several 10-mL portions of cold water. Rinse the

outside of the crucibles to remove residual (NH 4 ) 2 C 2 O 4 , and return the precipitate to the beakers in which

the CaC 2 O 4 was formed. Precipitation will occur as the following reaction:

Ca

(aq) +^ C 2 O 4

(aq) +^ H 2 O(l) ^ CaC 2 O 4 .H 2 O(s)

Titration

Add 100 mL of water and 50 mL of 3 M H 2 SO 4 to each of the beakers containing the precipitated

calcium oxalate. Heat to 80°C to 90°C, and titrate with 0.02 M permanganate. The temperature should be

greater than 60°C throughout the titration; reheat if necessary. Report the percentage of CaO in the

unknown. Redox titration results this reaction:

5 C 2 O 4

• 2 MnO 4

• 16 H

 10 CO 2 + 2 Mn

• 8 H 2 O

Note

The period of standing can be longer if the unknown contains no Mg 2+ .

Experiment 4. Determination of Hardness in Water Samples with EDTA

PRINCIPLE:

Water hardness, due to Ca

2+ and Mg

2+ , is expressed as mg/l CaCO 3 (ppm). The total of Ca

2+

and Mg

2+ is titrated with standard EDTA using ErioChrome Black T indicator. A standard EDTA

solution is prepared from dried Na 2 H 2 Y.2H 2 O. If the sample does not contain magnesium, Mg-

EDTA is added to titration flask to provide a sharp end point with the ErioChrome Black T, since

calcium does not form a sufficiently strong chelate with the indicator to give a sharp end point.

Equations

Titration:

End Point:

Ca

2+

• H 2 Y

2 - CaY

2 -

• 2H

Ca 2+

• MgY 2 - CaY 2 -
• Mg 2+

Mg

2+

• Hln

2 - Mgln

• • H

Mgln

• • H 2 Y

2 - MgY

2 -

• Hln

2 -

• H

(red) (colorless) (colorless) (blue)

The free acid parent of indicator is H 3 In, and that of the titrant EDTA is H 4 Y.

SOLUTIONS AND CHEMICALS REQUIRED:

1. ErioChrome Black T indicator 2. To prepare:

a) NH 3 - NH 4

buffer solution, pH 10. Dissolve 3.2 g NH 4 Cl in water, add 29 ml conc. NH 3 , and

dilute to about 50 ml. The buffer solution is best stored for a long period of time in a polyethylene

bottle to prevent leaching of metal ions from glass.

b) Standard 500 ml of 0.01 M EDTA solution. Dry about 3g reagent-grade Na 2 H 2 Y.2H 2 O in a

weighing bottle at 80

0 C 2 hr. Cool in desiccator and weigh necessary amount of Na 2 H 2 Y.2H 2 O.

PROCEDURE:

a) Standardization of 0.01 M EDTA

Pipet 10 mL standard 0.010 M Ca

solution into a clean 250 ml Erlenmeyer flask. Add 3 mL of

NH 3 - NH 4

buffer solution and dilute with distilled water. Then add small amounts of ErioChrome

Black T. Titrate Ca

solution with 0.01 M EDTA until color changes from wine red through

purple to a pure blue. Titrate at least three samples of standard Ca

solution with EDTA. Calculate

the concentrations of EDTA.

Experiment 5. Potentiometric Titration of Phosphoric Acid

PRINCIPLE:

In this experiment, a solution containing phosphoric acid H 3 PO 4 , will be titrated with

standardized NaOH solution. Measuring the pH of the solution after each addition of titrant will

monitor the titration. Analysis of the resulting titration curve will permit calculation of the exact

molarity of the phosphoric acid solution. The first and second acid dissociation equilibrium

constants for phosphoric acid will also be determined.

Phosphoric acid is a weak polyprotic acid that can dissociate stepwise as shown in

equations (1), (2), and (3). (H 3 PO 4  H 3 A)

H 3 A + H 2 O  H 2 A

• • H 30

…………………(1) Ka

H 2 A

• • H 2 O  HA

2 -

• H 30

…………………(2) Ka

HA

2 -

• H 2 O  A

3 -

• H 30

…………………(3) Ka

As the acid is titrated with strong base, the pH changes in a characteristic way giving rise to

a 2 step titration curve (do not observe third ionization constant (Ka3) on the titration curve). The

rate of change of H

ion concentration increases until it reaches a maximum rate at the equivalence

point.

SOLUTIONS AND CHEMICALS REQUIRED:

1. 250 mL of 0.1 M H 3 PO 4 and 250 mL of 0.1 M NaOH, 2. Standard buffer solutions at pH 4,7,10, for pH meter.

PROCEDURE:

1. Calibrate the pH meter with standard buffer solutions. 2. Transfer 10 mL of 0.1 M H 3 PO 4 into a 250 mL beaker and then add 90 mL of pure^ water. 3. Immerse the electrodes into solution after making sure the stirrer cannot hit the electrodes.

Measure pH of the solution and then titrate with sodium hydroxide solution with 0.5 mL

intervals but 0.1 mL intervals around the expected end points. Continue titrating until the pH is

about 11.

4. Plot the pH against the volume of the reagent.

Find the values of Ka1, Ka2 from the graph.

Experiment 6. Determination of Equilibrium Constants for Complex

Ions of Silver

PRINCIPLE:

The determination of the equilibrium constants often involve a titration, followed by

graphical analysis. A more direct approach for determining such values utilizes a concentration

cell and the Nernst equation. The experiment requires only simple equipment and in expensive

chemicals, and excellent results can be obtained in a relatively short time.

This method can be extended to include equilibria involving a host of metal ligand

complexes, as well as the determination of Ksp values for relatively insoluble salts.

The formation of a metal complex in an aqueous solution can be written as:

M

x+ (aq) + nL

y- (aq)  M Ln

( x+)+n(y-) (aq)

where the overall formation constant is

Kform = [M Ln

( x+)+n(y-) (aq)] / [M

x+ (aq)] [L

y- (aq)]

n

In the specific example described below, where x = 1 and y = 0, these equations can be abbreviated

as

and

Ag

• 2 NH 3  Ag(NH 3 ) 2

Kform = [Ag(NH 3 ) 2

] / [Ag

] [NH 3 ]

2

Knowledge of the three equilibrium concentrations clearly yields Kform

Several assumptions are made while doing this experiment:

1. The coordination number of metal ion and the formula of the complex are known. 2. The reaction reaches equilibrium quickly, 3. All activity coefficients are assumed to be 1.00 and junction potentials are negligible; 4. Excess amount of the ligand is added to the metal ion; thus, its concentration is decreased in

complex formation; but it can be neglected.

5. The overall formation constant, Kform, has a relatively large value, so the equilibrium

concentration of Ag(NH 3 ) 2

ion is practically the same as the concentration of initial Ag

ion.

6. Concentrations of intermediate species (those containing both water and the new ligand in the

coordination sphere) are negligibly small.

Because the ligand is in large excess, the equilibrium concentration of the newly formed

complex can be calculated by dividing millimoles of Ag

to the final volume. The concentration of

the remaining free ligand can then be determined from the balanced chemical equation.

The low equilibrium concentration of the aquated metal ion is the only missing datum to

determine Kform, and it can be obtained using a concentration cell and the Nernst equation. The emf

recorded. (The electrodes should not be in contact with the paper salt bridge.) In accord with

equation 1, students find that the voltmeter reads nearly zero.

2. (^) The solution in one of the beakers is poured into waste and a new half-cell is prepared. 5.0 mL

portion of AgNO 3 is transferred to a 150 mL beaker and then diluted to 50 mL with distilled water.

The Ag wire and salt bridge are reassembled as above and an additional reading is made, which

should be about 59 mV. Students are asked to define the cathode and anode and to rationalize the

electrode polarities with those of the voltmeter.

3. (^) One additional dilution is made, and the voltage is measured as described. The result is

compared with the calculated value.

The Kform for the diammine silver (I) complex is determined by adding a large but known

excess of ammonia to a known amount of aqueous AgNO 3. This is done as follows. The reference

half-cell is left undisturbed. Into a clean, dry a 50 mL beaker is added 15.0 mL of 0.100 M aqueous

ammonia and 15.0 mL of 0.010 M AgNO 3. The salt bridge is reassembled and the voltage

generated is measured by immersing Ag wire connected to voltmeter. Then concentration of the

Ag

(aq) ion remaining in the beaker is determined from the Nernst equation. The millimoles of the

complex formed and the free ammonia remaining, along with the final volume of the solution,

yield the molarities of these two species at equilibrium.

Students are asked to repeat the above procedure with different volumes of the Ag

and

NH 3 solutions and obtain a second value of this formation constant.

Experiment 7. Potentiometric Titration of Chloride and Iodide Mixtures

PRINCIPLE:

The mixture is titrated with a standard solution of silver nitrate, and the potentiometric end

points are indicated with a standard silver-wire electrode – glass electrode pair using a pH-meter

for potential measurements. Because the pH during the titration remains essentially constant, the

glass electrode’s potential remains constant, and this electrode serves as the reference electrode.

Thus, this eliminates the necessity of preparing a chloride-free salt bridge for the reference

electrode. AgI (Ksp= 1x

• 16 ) precipitates first since it is less soluble than AgCl (Ksp = 1x - 10 ).

The AgCl starts precipitating near the equivalence point of iodide titration

• • – 10 + – (^8) 1/2 – 8

(when [Ag ][Cl ] = 1x10 [Ag ] at the iodide equivalence point is (1x 10 ) =1x10 M). The

potential increment of the iodide titration curve will level off at the point when the chloride starts

precipitating, that is, near the iodide equivalence point inflection. This will be followed by the

typical S-shaped chloride potentiometric end point. The error in determining the iodide end point is

small if it is taken at the point at which the potential levels off.

I + Ag ^ AgI

Cl + Ag ^ AgCl

SOLUTIONS AND CHEMICALS REQUIRED:

1. 0.1 M standard AgNO 3 : Dry the primary standard AgNO 3 for 1-2 hours at 110- 120 0 C (no

longer). Store in a dessicator until it is ready for weighing. Obtain and dry your unknown at 120

o C

for 1-2 hours. Store in dessicator until it is ready for weighing.

PROCEDURE:

Obtain your unknowns from the instructor and dilute to approximately 150 mL and then

put a magnetic stirring bar, and place the beaker on a magnetic stirrer. Immerse the electrodes in

the solution, taking care that they do not hit the magnetic stirrer. Connect the silver electrode to the

reference terminal of the pH/ion meter and glass electrode to its usual terminal. Stir the solution

and titrate the sample with the standard AgNO 3. Take “pH” readings (actually pX) at 0.5 mL

increments until a significant increase is observed and then add 0.1 mL increments. After the first

end point is reached, add 0.5 mL increments until the second end point is approached and then

0.1 mL increments. Plot the potential versus volume of AgNO 3 and determine the end point for the

iodide and the chloride. Use these values to estimate the end point for the other two samples and