NORTHERN CARIBBEAN UNIVERSITY

DEPARTMENT of BIOLOGY, CHEMISTRY and

ALLIED HEALTH SCIENCES

LABORATORY MANUAL

CHEM202: ANALYTICAL CHEMISTRY II

Prepared by: Dr Nicole White

TABLE OF CONTENTS

PAGE

TABLE OF CONTENTS 1

CHEMISTRY LABORATORY REGULATIONS AND SAFETY

PRECAUTIONS 2

EXPERIMENT 1: THE DETERMINATION OF ACID CONTENT

IN VINEGARS 8

EXPERIMENT 2: THIN LAYER CHROMATOGRAPHY 9

EXPERIMENT 3: PREPARATION OF A STANDARD SILVER

NITRATE SOLUTION 13

EXPERIMENT 4: DETERMINATION OF CHLORIDE CONTENT

BY THE MOHR METHOD 14

EXPERIMENT 5: CHLORIDE CONTENT IN DIFFERENT

SOURCES OF WATER 16

EXPERIMENT 6: GRAVIMETRIC DETERMINATION OF CHLORIDE

IN A SOLUBLE SAMPLE 18

EXPERIMENT 7: DETERMINATION OF CALCIUM BY

DISPLACEMENT TITRATION 20

EXPERIEMENT 8: DETERMINING COPPER CONTENT IN SOIL

USING ATOMIC ABSORPTION SPECTROSCOPY 22

CHEMISTRY LABORATORY REGULATIONS AND SAFETY PRECAUTIONS

Compiled by Nicole White (PhD) 2008

1

Reference: http://www.sciencebyjones.com/safety_rules.htm

Thanks to the Flinn Scientific Safety Rules for much of the above.

  General Guidelines

 1. Conduct yourself in a responsible manner at all times in the laboratory.

2. Be familiar with your lab assignment before you come to lab.  Follow all written and verbal instructions carefully. If you do not understand a direction or part of a procedure, ask the teacher before proceeding.

3. Never work alone. No student may work in the laboratory without an instructor present.

4. When first entering a science room, do not touch any equipment, chemicals, or other materials in the laboratory area until you are instructed to do so.

5. Do not eat food, drink beverages, or chew gum in the laboratory. Do not use laboratory glassware as containers for food or beverages.

6. Perform only those experiments authorized by the instructor. Never do anything in the laboratory that is not called for in the laboratory procedures or by your instructor. Carefully follow all instructions, both written and oral. Unauthorized experiments are prohibited.  

7. Safety goggles and aprons must be worn whenever you work in lab.  Gloves should be worn whenever you use chemicals that cause skin irritations or need to handle hot equipment.  Wear older clothes that cover the maximum amount of skin. 

8. Observe good housekeeping practices. Work areas should be kept clean and tidy at all times. Bring only your laboratory instructions, worksheets, and/or reports to the work area. Other materials (books, purses, backpacks, etc.) should be stored in the classroom area.

9. Know the locations and operating procedures of all safety equipment including the first aid kit, eyewash station, safety shower, spill kit, fire extinguisher, and fire blanket. Know where the fire alarm and the exits are located. 

10. Be alert and proceed with caution at all times in the laboratory. Notify the instructor immediately of any unsafe conditions you observe.

11. Dispose of all chemical waste properly. Never mix chemicals in sink drains. Sinks are to be used only for water and those solutions designated by the instructor. Solid chemicals, metals, matches, filter paper, and all other insoluble materials are to be disposed of in the proper waste containers, not in the sink. Check the label of all waste containers twice

Compiled by Nicole White (PhD) 2008

2

before adding your chemical waste to the container.  Cracked or broken glass should be placed in the special container for “Broken Glass.” 

12. Labels and equipment instructions must be read carefully before use.  Set up and use the prescribed apparatus as directed in the laboratory instructions provided by your teacher.

13. Keep hands away from your face, eyes, mouth, and body while using chemicals.  Wash your hands with soap and water after performing all experiments.  Clean (with detergent powder), rinse, and dry all work surfaces and equipment at the end of the experiment. 

14. Experiments must be personally monitored at all times. You will be assigned a laboratory station at which to work. Do not wander around the room, distract other students, or interfere with the laboratory experiments of others.

15. Students are never permitted in the science storage rooms or preparation areas unless given specific permission by their instructor.

16. Know what to do if there is a fire drill during a laboratory period; containers must be closed, gas valves turned off, fume hoods turned off, and any electrical equipment turned off.

17. If you spill acid or any other corrosive chemical on you skin or clothes immediately wash area with large amounts of water (remember that small amounts of water may be worse that no water at all).  After this get the teacher’s attention.  The spill kit will be used for spills on floor or counter-top.

18. At the end of the laboratory session see that:  a) main gas outlet valve is shut off   b) the water is turned off   c) desk top, floor area, and sink are clean    d) all equipment is cool, clean, and arranged.

 Clothing

19. Any time chemicals, heat, or glassware are used, students will wear laboratory goggles. There will be no exceptions to this rule!  Contact lenses should not be worn in the laboratory unless you have permission from your instructor.

20. Dress properly during a laboratory activity. Long hair, dangling jewelry, and loose or baggy clothing are a hazard in the laboratory. Long hair must be tied back and dangling jewelry and loose or baggy clothing must be secured. Shoes must completely cover the foot. No sandals are allowed.

  

Compiled by Nicole White (PhD) 2008

3

Accidents and Injuries

21. Report any accident (spill, breakage, etc.) or injury (cut, burn, etc.) to the instructor immediately, no matter how trivial it may appear.

22. If you or your lab partner is hurt, immediately yell to get the instructor's attention. Everyone should turn off burners and prepare to help if needed.

23. If a chemical should splash in your eye(s), immediately flush with running water from the eyewash station for at least 20 minutes. Notify the instructor immediately.

 Handling Chemicals

24. All chemicals in the laboratory are to be considered dangerous.  Do not touch, taste, or smell any chemical unless specifically instructed to do so. The proper technique for smelling chemical fumes (when instructed to do so by the teacher) is to gently fan the air above the chemical toward your face.  Breathe normally.

25. Check the label on chemical bottles twice before removing any of the contents. Take only as much chemical as you need.  Smaller amounts often work better than larger amounts.   Label all containers and massing papers holding dry chemicals.

26. Never return unused chemicals to their original containers.

27. Never use mouth suction to fill a pipette. Use a pipette bulb or pipette filler.

28. Acids must be handled with extreme care.  ALWAYS ADD ACID SLOWLY TO WATER, with slow stirring and swirling, being careful of the heat produced, particularly with sulfuric acid.

29. Handle flammable hazardous liquids over a pan to contain spills. Never dispense flammable liquids anywhere near an open flame or source of heat.

30. Never take chemicals or other materials from the laboratory area.

31. Take great care when transferring acids and other chemicals from one part of the laboratory to another. Hold them securely and in the method demonstrated by the teacher as you walk.

 Handling Glassware and Equipment

32. Inserting and removing glass tubing from rubber stoppers can be dangerous. Always lubricate glassware (tubing, thistle tubes, thermometers, etc.) before attempting to insert it in a stopper. Always protect your hands with towels or cotton gloves when inserting glass tubing into, or removing it from, a rubber stopper. If a piece of glassware becomes "frozen" in a stopper, take it to your instructor for removal.

Compiled by Nicole White (PhD) 2008

4

33. When removing an electrical plug from its socket, grasp the plug, not the electrical cord. Hands must be completely dry before touching an electrical switch, plug, or outlet.

34. Examine glassware before each use. Never use chipped or cracked glassware. Never use dirty glassware.   Do not immerse hot glassware in cold water; it may shatter.

35. Report damaged electrical equipment immediately. Look for things such as frayed cords, exposed wires, and loose connections. Do not use damaged electrical equipment.

36. If you do not understand how to use a piece of equipment, ask the instructor for help.

 

Heating Substances

37. SHOULD THE BUNSEN BURNER GO OUT, IMMEDIATELY TURN OFF THE GAS AT THE GAS OUTLET VALVE.  If you wish to turn off the burner, do so by turning off the gas at the gas outlet valve first, then close the needle valve and barrel. Never reach over an exposed flame. Light gas burners only as instructed by the teacher.

38. Never leave a lit burner unattended. Never leave anything that is being heated or is visibly reacting unattended. Always turn the burner or hot plate off when not in use.

39. You will be instructed in the proper method of heating and boiling liquids in test tubes. Do not point the open end of a test tube being heated at yourself or anyone else.

40. Heated metals, glass, and ceramics remain very hot for a long time. They should be set aside to cool on a wire gauze and then picked up with caution. Use tongs or heat-protective gloves if necessary. Determine if an object is hot by bringing the back of your hand close to it prior to grasping it.

 

 

Compiled by Nicole White (PhD) 2008

5

GENERAL NOTES The Abstract should provide a less-than-150-word summary of the entire work: the

purpose, procedure, key results, and their significance should all be briefly addressed in this essential part of your report. The Abstract is not the place to introduce the experiment or describe the underlying principles in any detail. Stated in another way, the paper really begins with the Introduction, not the Abstract. Most scientists write the Abstract after they have written the rest of the paper, since it summarizes the work described. Never present material in the Abstract that you have not also presented somewhere in the main body of the report.

The Introduction should describe the specific goals of your experiment. What have youanalyzed, and why? Briefly discuss some of the key ideas. You are not required to use other references, but you are welcome to.

The Procedure should provide a concise description of how the experiment was actuallyconducted. Note important observations (especially events that likely introduced error) and highlight any deviations from the instructions in the handout you do not need to include drawings of any apparatus used in the experiment unless you feel it will aid your discussion.

Calculating acid concentrations from reagent stock

The molarity (cm) and percentage (cp) relationship depends on the density of solution (d) along with the molecular mass (MM) of the dissolved substance. Two equations depicting the interconversion of these two are as follows:

cm = cp × d / (100% × MM)

Reference: http://www.trimen.pl/witek/calculators/stezenia.html http://www.macalester.edu/~kuwata/classes/2005-06/Chem%20222/Pb%20in%20Soil

%20Lab%202006.pdf

Compiled by Nicole White (PhD) 2008

6

GENERAL LAB EQUIPMENT

Reference1. http://images.google.com/imgres?imgurl=http://physics.gallaudet.edu/classes/chemappa.gif&imgrefurl=https://sites.google.com/a/ggwo.org/ggca_science_lab/Home/chemistry/chemistry-equipment--safety&h=851&w=680&sz=140&hl=en&start=2&um=1&usg=__UvvQ2Dnd800KdbfMqNavZKP_T_Y=&tbnid=KsuthEI8w4YH9M:&tbnh=145&tbnw=116&prev=/images%3Fq%3Dlab%2Bequipment:%2Bchemistry%26um%3D1%26hl%3Den%26rls%3Dcom.microsoft:en-gb:IE-SearchBox%26rlz%3D1I7RNWM%26sa%3DN2.http://images.google.com/imgres?imgurl=http://visual.merriam-webster.com/images/science/chemistry/laboratory-equipment_2.jpg&imgrefurl=http://visual.merriam-webster.com/science/chemistry/laboratoryequipment_2.php&h=384&w=550&sz=55&hl=en&start=3&um=1&usg=__3DXNsK64lkexY4xcVrs89SaC0PU=&tbnid=VenjDu5mM2Rq5M:&tbnh=93&tbnw=133&prev=/images%3Fq%3Dlab%2Bequipment:%2Bchemistry%26um%3D1%26hl%3Den%26rls%3Dcom.microsoft:en-gb:IE-SearchBox%26rlz%3D1I7RNWM%26sa%3DN3. http://visual.merriam-webster.com/images/science/chemistry/laboratory-equipment_1.jpg 4. http://bioweb.wku.edu/courses/Biol121/Carbo/pipets.png5. http://www.naugraexport.com/glass/front.jpg 6. http://www.science-laboratory-equipment.com/product-images/1059.jpg7. http://www.unisciencelab.com/unisciencelab/laboratory/images/tongs_1.jpg 8. http://z.about.com/d/chemistry/1/0/D/U/potferri.jpg

Compiled by Nicole White (PhD) 2008

7

EXPERIMENT 1

THE DETERMINATION OF ACID CONTENT IN VINEGARS

Aim To determine the acid contents in various brands of vinegars.

IntroductionVinegar or French for sour wine is formed by aerobic bacteria oxidizing grain alcohol to acetic acid and water. More generally, vinegar can be defined as a solution composed of acetic acid (HC2H3O2), water, and, perhaps, other substances. To be sold in stores as vinegar, this solution must contain at least four grams of acetic acid per 100 ml of solution. For the titration of the vinegar in this experiment the following specific reaction will be used to calculate the acetic acid content of the vinegar sample:H3C2O2H(aq) + NaOH(aq) → H2O(l) + C2H3O2Na (aq).

% w/v = (mass of solute/volume of solution)×100%.

ProcedureKeep flask closed when not in use as the acidity of bottled vinegar tends to decrease on exposure to air. Record the brand of vinegar used.

1. Standardize the ~ 0.1NaOH solution using potassium hydrogen phthalate.2. Pipette 25 mL of the unknown into a 250 mL volumetric flask and dilute to the mark with

distilled water.3. Pipette 50 mL aliquot of vinegar into a 250 mL conical flask. Add 50 mL of water.4. Titrate against the NaOH solution.

Titrations are to be done in triplicate.

Treatment of data Determine the concentration of NaOH used. Determine the w/v% vinegar in each sample.

Questions1. Does the vinegar studied meet the commercial law specification of a minimum of 4 g of

acetic acid/100 ml of vinegar? 2. Has the vinegar supplier truthfully reported the percent acidity? Why is it a good idea to

rinse the burette with the NaOH solution, instead of with water, before filling it at the start of the titration?

Reference http://www.baruch.cuny.edu/wsas/departments/natural_science/chemistry/chm_1000/vinegar.pdf

Skoog, D.A.; west, D.M.; Holler, F.J. Foundations of analytical Chemistry, 7th Ed.; Saunders College Publishing: 1991

Compiled by Nicole White (PhD) 2008

8

EXPERIEMENT 2

THIN LAYER CHROMATOGRAPHY

Aim To investigate the separation of mixtures with thin-layer-chromatography and calculate

Rf values To determine the composition of various over-the-counter analgesics

IntroductionSince the investigation of chemically pure materials has provided the basis for understanding chemistry, it is critical that chemists be able to separate mixtures of substances as they are often found in nature. Thin layer chromatography is a special application of adsorption in which a thin layer of adsorbent supported on a flat surface is utilized instead of a column of adsorbent. The most commonly used adsorbent is silica gel and the flat surface is a plain glass plate (a cut piece of sheet glass). All forms of chromatography depend upon the distribution (or separation) of solute particles between a moving phase (a gas or liquid; you'll be using a liquid moving phase in this procedure) and the stationary phase (a liquid or solid; in this procedure you will be using a solid stationary phase.The separation of the components of a mixture depends on adsorption-desorption equilibria between compounds adsorbed on the solid stationary phase and in the moving liquid phase. The extent of adsorption of a single component depends upon the polarity of the molecule, the activity of the adsorbent, and the polarity of the mobile liquid phase. The actual separation of the components in a mixture is dependent on the relative values of the adsorption-desorption equilibrium constants for each of the components in the mixture. In general, the more polar a functional group in the compound is, the more strongly it will be adsorbed on the surface of the solid phase. If the components are colored, they can be located readily, but more often they are invisible and must be located by other means. Illumination with ultra-violet light will excite fluorescence in many compounds.

Because each chromatogram will be developed under slightly different conditions (slightly more or less solvent, slightly different temperatures, differing size and concentration of the origin line applications), the general practice in TLC is to place known and unknown compounds on the

Compiled by Nicole White (PhD) 2008

9

same chromatogram. The unknowns can then be easily identified by comparison with the known samples that have been run under the exact same conditions on the same chromatogram. A few characteristics that remain consistent from run to run can be used to compare different chromatograms. The generally accepted method for making such comparisons is to compute the Rf value (meaning "Relationship to the Front") of the spots. This is easily calculated by:

Rf = Distance moved by compound Distance moved by solvent system

(Note that a more polar compound will more strongly adsorb to the polar stationary phase and will have a lower Rf value)

The means, therefore, available to identify the components of a TLC mixture are:• Migration identical to that of a standard or reference (known) compound• Intrinsic fluorescence like that of a reference compound• Staining behavior like that of a reference compound (Staining characteristics may be concentration-related. Therefore, a spot of lower concentration may have a lighter color than a spot of higher concentration) • The Rf value that is identical to that of a reference compound. This Rf value may be obtained from direct measurement or obtained from literature.

You will be using TLC to separate and identify the components of some familiar mixtures of organic compounds – over-the-counter analgesic drugs. Several of these compounds have related structures, but are different enough so that they can be separated.Aspirin is acetylsalicylic acid – made from the esterification reaction of salicylic acid with acetic anhydride. It was first used in ancient times, when the Greek physician Hippocrates had his patients chew willow bark to ease pain and reduce fever. It was later discovered that willow bark contains salicin, a derivative of the molecule we know as aspirin today. Aspirin is an analgesic (pain reliever) as well as an antipyretic (fever reducer). Its action in both cases, analgesic and antipyretic, is due to the ability of aspirin to inhibit prostaglandins, which are chemicals in the body involved in sending pain messages to the brain. Any analgesic doesn't really stop the ache or pain; an analgesic just keeps the brain from hearing about the pain. Aspirin also can inhibit theclotting of blood, so sometimes it is prescribed for lowering the risk of heart attack and stroke (cerebral vascular accidents) by lowering the formation of blood clots that can dislodge and plug arteries or veins in the heart and brain.Acetaminophen is structurally related to aspirin, and it is an analgesic and an antipyretic, but does not reduce inflammation like aspirin does. It also does not promote gastrointestinal upset like aspirin does, so it can be taken in larger doses over longer periods of time.Caffeine is often found in preparations of over-the-counter analgesics, but there is little evidence that caffeine enhances the effects of aspirin. Caffeine is a mild stimulant. Salicylamide is also structurally related to the aspirin, and is sometimes used in combination with aspirin.Ibuprofen is also an analgesic (but not used in this procedure). It only bears a slight structural resemblance to aspirin, but may be superior to aspirin in the reduction of inflammation. It can also reduce fever and relieves mild pain, but does not appear to be better than aspirin for these effects.

Compiled by Nicole White (PhD) 2008

10

ProcedureUse the ethyl acetate + 0.5% acetic acid solvent for this section. Be careful; don't shine the UV light in your eyes.

1. Obtain on thin layer chromatography plate from your instructor. Handle the plate by the edges only. With a pencil and ruler, lightly mark a line about 1 cm from the narrow edge of the plate (not the wide edge of the plate).

2. Divide the plate into six equal portions. Six solutions will be spotted on each plate (keeping the spots as small as possible, approximately 1-2 mm in diameter) – a solution of each of four standards/reference solutions (acetaminophen, aspirin, caffeine, salicylamide) will be applied to the plate, as well as a combined solution of REV 01/06 11 all four standards/references together, so it can be determined how the compounds run when together in solution. You will also spot your assigned unknown on the plate. Apply each of the six solutions once; then spot your unknown a second time, allowing the spot to dry before placing the other spot on top of it. It might be easiest to apply the solutions in alphabetical order, to prevent confusion. Apply each substance twice to plate on the same spot.

3. Dry the plates in an 80 °C oven for 2-3 minutes (especially if the weather is humid). The chromatograms will be developed in a mixture of ethyl acetate/acetic acid.

4. While the plates are drying in the oven, prepare the chromatography chamber. You will use your 400-mL beaker as a development chamber. Place enough of the solvent (ethylacetate/acetic acid) into the bottom of the chromatography so that the depth of solvent is enough to cover the bottom of the chamber, but not enough to touch the applied spots.

5. Insert the spotted, dried plate into the prepared chamber and cover the chamber with a plastic sandwich bag. Do not move the jar during development. Watch the progress of the solvent front on the plate. When the solvent front is about 0.5-1.0 cm from the top of the plate, remove the plate from the chamber and QUICKLY mark the solvent front level with a pencil line across the top of the chromatogram before the solvent evaporates. After marking the solvent front, allow the chromatogram to dry. Place the used solvent in the proper waste container. Clean the chamber (a 400-mL beaker) with soap and water. Rinse the beaker with deionized water as a final rinse.

6. Visualize the plate under short-wave UV light, and mark any spots (either fluorescent or dark spots) that are seen with a pencil on your chromatogram. Then place the chromatogram into a jar with a few iodine crystals, cover the jar and let stand until brownish-yellow spots develop. Remove the plates and outline the spots with a pencil.

7. Notice the placement of the compounds:Analgesic Rf Value

Acetaminophen (A) 0.476Salicylamide (S) 0.738

Aspirin (Asp)0.3620.619

Ibuprofen (I) 0.785Caffeine (C) 0.362

Compiled by Nicole White (PhD) 2008

11

Aspirin and salicylamide run at about the same Rf value; the way they are told apart is by the visualization process. If both compounds are present, both visualization steps will show the presence of the proper compounds.

Treatment of data Make a drawing of this chromatogram in your laboratory notebook, indicating the proper

colour and shape of each spot. Measure the distance from the origin line to the solvent front. Measure all distances in

centimeters and measure to the nearest 0.01 cm (two decimal places). Then measure the distance from the origin line to the center of each spot (again, to the nearest 0.01 cm). Write these distances on the drawing in your notebook.

Using these measured distances, calculate the Rf of these spots. Determine the components

Questions1. What would happen if the solvent layer in your development chamber is too deep, so that

the origin lines of your chromatograms are submerged in it?2. Using an appropriate source, draw the structural formula of the following compounds in

the spaces provided above each name. a) aspirin b) ibuprofen c) acetaminophen d) caffeine

Reference- http://www.ipfw.edu/chem/112/kimble/2-Thin%20Layer%20Chromatography.pdf

- http://delloyd.50megs.com/labscripts/TLC.html

- http://bama.ua.edu/~kshaughn/ch338/handouts/TLC-exp1.pdf

Compiled by Nicole White (PhD) 2008

12

EXPERIMENT 3

PREPARATION OF A STANDARD SILVER NITRATE SOLUTION

Aim To utilize proper analytical techniques in preparation of a standard solution

IntroductionMost precipitation titrations make use of a silver nitrate solution as titrant. It offers several other advantages as well. It is non-hygroscopic, in contrast to silver fluoroborate and silver perchlorate. It is relatively stable to light. Finally it dissolves in numerous solvents. The nitrate can be easily replaced by other ligands, rendering AgNO3 versatile. Treatment with solutions of halide ions gives a precipitate of AgX (X = Cl, Br, I).

CAUTIONAs with all silver salts, silver nitrate is toxic and corrosive. Brief exposure to the chemical will not produce immediate or even any side effects other than the purple skin stains, but with more exposure, side effects will become more noticeable. It is also very poisonous and can cause burns. Long-term exposure may cause eye damage. Short contact can lead to deposition of black silver stains on the skin. Besides being very destructive of mucous membranes, it is a skin and eye irritant.

Procedure1. Weigh approximately 16.9 g of AgNO3 and dry at 110 oC for no more than an hour.

Prolonged heating causes partial decomposition of AgNO3. Cool to room temperature in a dessicator.

2. Weigh the bottle and contents to the nearest 0.0001 g.3. Transfer the bulk to a 1000 mL volumetric flask using a funnel.4. Cap the weighing bottle and re-weigh.5. Rinse the funnel thoroughly.6. Dissolve the AgNO3. Fill flask to the mark.

Treatment of data Calculate the molar concentration of this solution.

Reference http://en.wikipedia.org/wiki/Silver_nitrate#Preparation Skoog, D.A.; west, D.M.; Holler, F.J. Foundations of analytical Chemistry, 7th Ed.;

Saunders College Publishing: 1991

SAVE SOLUTION FOR NEXT 3 LABS. STORE IN A DARK PLACE.

Compiled by Nicole White (PhD) 2008

13

EXPERIMENT 4

DETERMINATION OF CHLORIDE CONTENT BY THE MOHR METHOD

Aim To find the percentage of Cl- ions in an unknown compound.

IntroductionBased on the chemical reaction, the chloride ions precipitated as silver chloride in the final titration experiment.

AgNO3 + NaCl → AgCl + NaNO3

The Mohr method uses the indicator potassium chromate. This method determines the chloride ion concentration of a solution by titration with silver nitrate. As the silver nitrate solution is slowly added, a precipitate of silver chloride forms.

Ag+(aq) + Cl–(aq) → AgCl(s)The end point of the titration occurs when all the chloride ions are precipitated. Then additional chloride ions react with the chromate ions of the indicator, potassium chromate, to form a red-brown precipitate of silver chromate.

2 Ag+(aq) + CrO42–(aq) → Ag2CrO4(s)

The Mohr titration should be carried out under conditions of pH 6.5 – 9. At higher pH silver ions may be removed by precipitation with hydroxide ions, and at low pH chromate ions may be removed by an acid-base reaction to form hydrogen chromate ions or dichromate ions, affecting the accuracy of the end point.

ProcedureDry reagent grade NaCl and unknown in oven for ca. 1 hr. at 110 oC. Cool in a dessicator.

1. Prepare 0.1 M AgNO3 solution. Express the concentration as weight molarity (mm AgNO3/g of solution)

2. Dissolve 1.0 g K2CrO4 in approximately 20 mL of distilled water.To standardize AgNO3:

3. Weigh 0.25 g portions of NaCl into 250 mL Erlenmeyer flask. Dissolve in 100 mL distilled water.

4. Add small quantities of NaHCO3 until effervence stops.5. Add ~2 mL K2CrO4 solution to NaCl solution. 6. Titrate against AgNO3 until you observe the first permanent red Ag2CrO4.7. Determine an indicator blank by suspending a small amount of chloride free CaCO3 in

100 mL of distilled water containing 2 mL of K2CrO4. Titrate as above.Unknown sample:

8. Collect unknown from instructor.9. Weigh accurately 0.2 g of sample. Titrate as above.

Treatment of data Correct reagent weights for the blank. Calculate %Cl- in NaCl.

Compiled by Nicole White (PhD) 2008

14

Determine % Cl- in unknown. Compare your result with the other groups.

Reference http://www.neiu.edu/~rdpurtuc/iSchool/chem/quant%20213/Quant-Lab-02.pdf Skoog, D.A.; west, D.M.; Holler, F.J. Foundations of analytical Chemistry, 7th Ed.;

Saunders College Publishing: 1991 http://www.mcm.edu/~simpsong/Quant2.htm http://www.chemteach.ac.nz/investigations/documents/chloride_mohr.pdf

Compiled by Nicole White (PhD) 2008

15

EXPERIMENT 5

CHLORIDE CONTENT IN DIFFERENT SOURCES OF WATER

Aim To determine the percentage of chlorine present in water from various sources

IntroductionTests for the presence of chloride ion are frequently required when testing water supplies for chlorination. In the pulp and paper industry, bleaching plants must do extensive analysis for chlorides before returning waste water to a river or stream. Treatment for chloride removal is costly and requires some level of technical expertise. In general, water having the highest quality is reserved for the potable water supply. It meets all international standards for drinking water and contains a chloride concentration less than 250 mg/L. Industrial water use varies widely (boiler-feed makeup, process water, truck washing, toilet flushing, etc) and therefore the water quality also varies widely. Agriculture can generally accept water having lower quality standards, but the salinity of irrigation water depends upon crop tolerance and rainfall. Irrigation water having more elevated levels of salinity can be tolerated by most crops in areas having high rainfall, evenly distributed throughout the year

 The procedure for testing for chloride ions in solution is fairly simple. Most chloride tests involve titration with silver nitrate solution in the presence of chromate ion which acts as an indicator. Chloride ions can be precipitated with silver nitrate to form insoluble silver chloride. 

Ag+(aq)) + Cl- (aq)) à AgCl (s) Silver chloride is less soluble than silver chromate. Therefore when the chloride has been removed from the solution the orange-red silver chromate begins to precipitate signaling the end point. 

2 Ag+(aq)) + CrO42- à Ag2CrO4 (s)

 In this experiment you will be determining the concentration of Chloride ion in seawater. The usual concentration is about 1.93% but the value varies slightly depending on the depth and the location where the sample was collected.

Compiled by Nicole White (PhD) 2008

16

ProcedureSilver nitrate will stain the skin and can be harmful if taken internally. Potassium Chromate may be a suspected carcinogen if inhaled or airborne. Avoid all contact with the skin

1. Weigh the pipette containing the salt water solution. Then put 20 drops of the salt water solution into a well of the 24 well plate or a small test tube. The well or test tube should be one quarter to one third full.

2. Carefully add 2-3 drops of the potassium chromate solution and stir with a toothpick.3. Record the initial mass of the silver nitrate pipette. 4. Add the silver nitrate a drop at a time while stirring with the toothpick or swirling the

solution between each additions. 5. Continue adding silver nitrate a drop at a time until a red precipitate persists throughout

the liquid and cannot be stirred away. 6. For sea water about 20-30 drops will be required. Reweigh the silver nitrate pipette to

determine the mass of AgNO3 added. 7. Repeat experiment for: pool water, tap water, and bottled water.

Treatment of data Tabulate you results as follows.

Mass/g Sample dataInitial mass salt water dropperSalt water dropper after adding water to test tubeMass of salt waterInitial mass of the dropper with silver nitrateMass of dropper after adding silver nitrate to test tubeMass of silver nitrate solution used

Calculate the mass of salt water used. Calculate the mass of silver nitrate used.   Find the number of moles of silver nitrate used. (Multiply the mass in calculation 2 by

the density of the solution and the concentration of silver nitrate solution in moles per dm3) 

5. Find the mass of Chloride ion present  6. Find the percentage of Chloride ion in solution

Reference www. sidsnet.org /docshare/other/20031105152734_CUBA.CASE_STUDY2.doc

http://lincoln.pps.k12.or.us/pages/lscheffler/ChlorideAnalysis.htm

PREPARE CRUCIBLES FOR NEXT WEEKS EXPERIMENT.

Compiled by Nicole White (PhD) 2008

17

EXPERIMENT 6

GRAVIMETRIC DETERMINATION OF CHLORIDE IN A SOLUBLE SAMPLE

Aim To accurately determine the mass of a substance by employing crucibles

IntroductionGravimetric analysis is a quantitative (i.e. how much?) method of classical analysis. The element to be determined is isolated in a solid compound of known identity and definite composition. The chloride content of a soluble salt can be determined by precipitation as silver chloride:

Ag+ + Cl- →AgCl(s)

In most cases, the analyte must first be converted to a solid by precipitation with an appropriate reagent. The precipitate can then be collected by filtration, washed, dried to remove traces of moisture from the solution, and weighed. The amount of analyte in the original sample can then be calculated from the mass of the precipitate and its chemical composition.

ProcedurePreparation of crucibles and preliminary lab work

1. Clean 2 medium porosity sintered glass by allowing them to stand in approximately 5 mL of concentrated nitric acid for 5 minutes.

2. Using vacuum filtration, draw the acid the crucible.3. Rinse thoroughly with tap water and finally with distilled water.4. Allow each crucible and dry to constant weight in an oven at 110 oC. Three constant

readings between 0.2 -0.3 mg are sufficient. Do not touch crucibles with hand once the drying process has begun.

5. Place crucibles in a dessicator until required for use. Allow it to cool to room temperature.

6. Transfer unknown to a 400 mL beaker. 7. Dissolve in 100 mL distilled water. Add 2-3 mL of HNO3.8. Slowly, and with good stirring, add 0.2 mL AgNO3 to each cold sample solutions until

AgCl begins to coagulate. Use separate stirring rods for each beaker.9. Add an adiitional 3 mL of AgNO3. Heat to almost boiling.10. Digest for 10 minutes. Add a few drops of AgNO3 to confirm precipitation is complete.11. If incomplete (more ppt. forms), then add 3 mL of AgNO3, digest, and again test for

precipitation.12. Cover each beaker and store in a dark place until the next laboratory session

Compiled by Nicole White (PhD) 2008

18

Experiment13. Decant the supernatant liquids through weighed filtering crucibles (previously prepared).

Wash the precipitates several times with a solution of 2-5 mL of 6 M HNO3 per litre of distilled water.

14. Filter, using vacuum filtration, liquids through crucibles. Use policemen to dislodge any particles that adhere to the walls of the beakers. Note which beaker was paired to which crucible.

15. Dry precipitates in an oven at 110 oC for at least 1 hr. Store the crucibles in a dessicator while they cool.

16. Determine the mass of crucibles and their contents. 17. Repeat the cycle of heating, cooling, and weighing (30 minute periods) until constant

weights are obtained.18. Place solids into paper and finally placed in waste provided. All solutions should also be

placed in waste containers.19. Clean crucibles by repeating steps 1-3.

Treatment of data Tabulate all masses, properly records must be kept. Calculate the percentage of Cl- in the sample.

Reference http://www.earlham.edu/~chem/chem111f03/labs/Experiment%206.pdf http://en.wikipedia.org/wiki/Gravimetry Skoog, D.A.; west, D.M.; Holler, F.J. Foundations of analytical Chemistry, 7th Ed.;

Saunders College Publishing: 1991

Compiled by Nicole White (PhD) 2008

19

EXPERIMENT 7

THE DETERMINATION OF CALCIUM BY DISPLACEMENT TITRATION

Aim To determine the metal content in a complex

IntroductionA solution of the magnesium/EDTA complex is useful for the titration of cations that form more stable complexes than the magnesium complex but for which no indicator is available. Ethylenediaminetetraacetic acid (EDTA) is a reagent that forms EDTA-metal complexes with many metal ions (but not with alkali metal ions such as Na+ and K+). The concentration of the magnesium solution is not important; all that is necessary is that the molar ratio between Mg 2+

and EDTA in the reagent be exactly unity. In alkaline conditions (pH >9) it forms stable complexes with the alkaline earth metal ions Ca2+ and Mg2+. The EDTA reagent can be used to measure the total quantity of dissolved Ca2+ or Mg2+ ions in a water sample.

ProcedurePreparation of Solutions

1. Buffer solution, pH 10: Dilute 75 mL of concentrated NH4OH and 10.5 g NH4Cl- in sufficient distilled water to give 150 mL of solution.

2. Eriochrome Black T indicator: Dissolve 250 mg of the solid in a solution containing 15 mL of triethaloamine and 5 mL absolute ethanol.

3. Prepare 6M and 0.1 M NaOH solutions in 250 mL volumetric flasks.4. Prepare 6 M HCl solution in a 100 mL volumetric flask.5. MgSO4/ EDTA solution: Dissolve 3.72 g EDTA in 50 mL of distilled water and add 2.46

g MgSO4.7H2O. Add a few drops of phenolphthalein indicator and titrate against 0.1 M NaOH until the white solution turns pink.

6. Blank solution: Mix 10.00 mL distilled water + 2 drops of Erichrome Black indicator + 2 drops of pH 10 buffer solution + 1.00 mL of MgEDTA solution, and titrate dropwise against 0.01M EDTA.

Standardization of 0.01M EDTA against Ca2+ solution of known concentration7. Weigh 0.9320g EDTA and mix with 250 mL of distilled water to give 0.01M EDTA.8. Weigh 0.2503g CaCO3 and mixed it with 5.00 mL 6M HCl + 50.00 mL H2O.9. Heat solution to boil for 3-4 minutes; let solution cool to room temperature.10. Titrate solution, carefully, against 6 M NaOH. 11. Dilute solution to 250 mL with water12. Pipette 10mL of new solution and mix it with 2 drops of Erichrome Black indicator + 2

drops of pH 10 buffer solution + 1.00 mL of MgEDTA solution.13. Titrate the 10.00 mL sample solution against stock 0.01M EDTA.14.Pipette 20.00 mL of CaCO3 for triplicate titrations against EDTA

Compiled by Nicole White (PhD) 2008

20

Determination of unknown Calcium by displacement titration15. Weigh 0.5g (0.4999g) of unknown sample and mix with 5.00 mL 6M HCl + 50.00 mL

water.16. Heat solution to boil for 3-4 minutes; let solution cool to room temperature.17. Titrate solution against 6 M NaOH. 18. Dilute solution to 250 mL with H2O.19. Pipette10.00 mL of new solution and mix it with 2 drops of Erichrome Black indicator+ 2

drops of pH 10 buffer solution + 1.00 mL of MgEDTA solution.20. Titrate, in triplicate, the 10.00 mL sample solution with stock standardized EDTA

solution

Treatment of data Calculate EDTA molarity. Calculate mass percentage of Ca2+ in unknown sample.

Reference http://www.neiu.edu/~rdpurtuc/iSchool/chem/quant%20213/Quant-Lab-05.pdf Skoog, D.A.; west, D.M.; Holler, F.J. Foundations of analytical Chemistry, 7th Ed.;

Saunders College Publishing: 1991

Compiled by Nicole White (PhD) 2008

21

EXPERIMENT 8

DETERMINING COPPER CONTENT IN SOIL USING ATOMIC ABSORBTION

SPECTROSCOPY

Aim To determine the copper concentration in soil To understand and apply a calibration curve

IntroductionAtomic absorption spectroscopy (AA or AAS) is a technique for determining the concentration of a particular metal element in a sample. It is one of the commonest instrumental methods for analyzing for metals and some metalloids. The technique makes use of absorption spectrometry to assess the concentration of an analyte in a sample and relies therefore heavily on Beer-Lambert Law. The analyte concentration is determined from the amount of absorption.

In their elemental form, metals will absorb ultraviolet light when they are excited by heat. Each metal has a characteristic wavelength that will be absorbed. The AAS instrument looks for a particular metal by focusing a beam of uv light at a specific wavelength through a flame and into a detector. The sample of interest is aspirated into the flame. If that metal is present in the sample, it will absorb some of the light, thus reducing its intensity. The instrument measures the change in intensity. A computer data system converts the change in intensity into an absorbance. As concentration goes up, absorbance goes up. The researcher can construct a calibration curve by running standards of various concentrations on the AAS and observing the absorbances.

Sample treatment1. Approximately, 5 g of each sample was placed in a 50 mL volumetric flask and the

container filled to the mark with distilled water. 2. The samples were then run using an atomic absorption spectrometer at a wavelength of

324.75 nm. The samples were run in triplicates.

ResultsAbsorbance readings for standard copper samples

[Cu]/ mg/L Mean Absorbance (A)0.000 0.01850.502 0.04000.993 0.06002.005 0.09503.999 0.1550

Compiled by Nicole White (PhD) 2008

22

Absorbance readings for unknown copper samplesSample A

1 0.04342 0.04703 0.04984 0.04915 0.04816 0.04727 0.06538 0.06179 0.0523

Treatment of data Determine the concentration of each of the unknown samples from the calibration curve.

The absorbances have not been corrected for the blank. Make the necessary corrections.

What are the concentrations of the original samples?

Questions1. Why were the measurements carried out in triplicates?2. What is the purpose of the blank run?3. If you were a field analyst asked to determine the average copper content in a particular

parish, outline the procedure you would use to carry out this investigation.

Reference http://en.wikipedia.org/wiki/Atomic_absorption_spectroscopy http://elchem.kaist.ac.kr/vt/chem-ed/spec/atomic/aa.htm http://www.shsu.edu/~chemistry/primers/AAS.html http://www.gmu.edu/departments/SRIF/tutorial/aas/aas.htm

Special thanks to Dr M. Wilson and S.I.R.I.

Compiled by Nicole White (PhD) 2008

23