Lab_Manual_STK1211_2015_2016

8/19/2019 Lab_Manual_STK1211_2015_2016

1/38

STK1211

Practical AnalyticalChemistry

Laboratory Manual

Department of ChemistryFaculty of Resource Science & Technology

Universiti Malaysia Sarawak2015/2016

8/19/2019 Lab_Manual_STK1211_2015_2016

2/38

2

Example of cover page

FACULTY OF RESOURCE SCIENCE AND TECHNOLOGYDEPARTMENT OF CHEMISTRY

STK 1211 – Practical Analytical Chemistry

EXPERIMENT NO:

TITLE OF EXPERIMENT:

DATE OF EXPERIMENT :

GROUP MEMBERS & MATRICNUMBERS

:

LAB FACILITATOR :

REPORT DUE DATE :

8/19/2019 Lab_Manual_STK1211_2015_2016

3/38

3

Table of Contents

Page

Laboratory outline

- Laboratory report format- Laboratory report submission- Group experiments

4

Laboratory safety 6

Apparatus and techniques

- “Clean” and “Clean and Dry” glassware - Desiccator and handling dried compounds- Electronic and analytical balance

- Volumetric flasks and quantitative transfers

7

Experiment 1 : Acid-base Titrations 11-14

Experiment 2 : Water Analysis: Suspended Solids and Dissolved Oxygen 15-19

Experiment 3 : Water Analysis: Hardness of Water and ChlorineConcentration

20-24

Experiment 4 : Redox Titration – Ascorbic Acid 25-29

Experiment 5 : Gravimetric Determination of Nickel(II) ion 30-33

Experiment 6 : Thin-layer Chromatography 34-38

8/19/2019 Lab_Manual_STK1211_2015_2016

4/38

4

Laboratory Outline

1. The lab report should be organized according to the following sequence:

a)

Introductionb) Objectivesc) Procedures and material/apparatusd) Resultse) Calculations, if anyf) Discussiong) Conclusionh) Post-lab questionsi) References

2. Unorganized lab report will be penalized .3. The content of each section in a lab report is described in Table 1.4. Only one lab report needs to be submitted per group. 5. Lab report is to be submit ted at the beginning of next lab session .6. Late submission will be penalized .7. Plagiarism is strictly prohibited .

Table 1: Description on the content of a lab report

Section DescriptionIntroduction Brief discussion on the background of the technique(s) used in the

experiment.The theory behind any calculation involved should be presented inthe introduction as well.

Any referred information must be cited.Objectives Brief description on the objectives of the outlined experimentExperimentalset up andprocedure

Consists of the list of chemicals and apparatus used. All steps performed in the experimental procedure should be listedin the order that they were performed, in exactly the manner inwhich you performed them.The procedures must be written in passive sentences.

Results Should list all data obtained, in raw form, with informationprovided as to how the data was obtained, as well as the accuracyof all measurements.Record the numbers/measurements you collect (Quantitative data)Record other pertinent observations (Qualitative data)

If possible use tables.Include graphs if appropriate.Include ALL, even those that will be rejected later.

Calculation(if any)

Calculation should be included after the results shown.Include formulas, units, use significant digits.Show complete calculations for all results.

8/19/2019 Lab_Manual_STK1211_2015_2016

5/38

5

Discussion The data should be discussed and evaluated, both positively andnegatively.Do not try to manipulate data to fit the results you think youshould obtain.Evaluate the data fairly, even if the data seem to contradict withtheory you may have been expecting the data to follow.

A discussion of possible sources of error should be included in thissection.Calculate your percent error, if applicable.If any data was rejected, explain why here.

Any referred information must be cited.Include chemical reactions involved (if any).

Conclusion Concise, direct statement of what you learned.If possible, use a single sentence.

Post-labquestions

Answer all the post-lab questions provided after each experiment.

References All materials have been used in writing the laboratory report

should be listed (at least two).Journal articles are referenced by listing the authors (last namefirst), the year of publication, the title of the article, the title of the

journal, the volume number and the page number. For example: Kuivinenm J., Johnsson, H. (1999) Determination of Trihalomethanesand some Chlorinated solvents in Drinking Water. Water Research, 33(5). 1201-1208.

Books are referenced by listing the authors (last name first), yearof publication, the title of the book, the edition, the publisher, cityof publication and the page number. For example:

John, M.M., Johnston, D.O., Nettervile, J.T., Wood, J.L., Joesten, M.D.

1991. Laboratory Manual to Accompany World of Chemistry. SoundersCollege Publishing, New York.

8/19/2019 Lab_Manual_STK1211_2015_2016

6/38

6

Laboratory Safety

Safety in the laboratory is a subject of the utmost importance. All chemicals are harmful tosome degree, therefore it is imperative to learn the safety rules and follow them strictly at alltimes. You will be expelled from the laboratory for failing to comply with these regulations.These rules are referred to many laboratories as “the usual safety procedures”.

General

1. Wear shoes at all times when you are in the laboratory.2. Wear lab coat at all times when you are in the laboratory3. Report any spill or accident immediately to your instructor.4. Know the location and operation of safety equipment in the laboratory from the first

meeting of the laboratory section.5. Drinking, eating and smoking are absolutely forbidden in the laboratory.6. Never work alone in the chemical laboratory.7. Dispose the chemicals properly, in the container provided, and according to the

instructions given by the laboratory instructor. Do not simply pour waste chemicalsdown the sink.

8. Keep your laboratory space clean.

Safety glasses

1. Safety glasses should be worn at all times while in the laboratory.2. Contact lenses should never be worn in the laboratory because they cannot be

removed rapidly if reagents accidentally splash in the eye.

Chemicals

1. Handle all chemicals according to any specific directions indicated on the container, orthose given to you by your instructor.

2. Avoid contact with skin and clothing.3. Wipe up spills immediately, especially near the balances and reagent shelf.4. Replace caps on containers immediately after use.5. Avoid the inhalation of organic vapours, particularly aromatic solvents and

chlorinated solvents.

Disposals of chemicals

1. Dispose of chemicals as directed in each experiment.2. Water-soluble substances can be flushed down the drain with large quantity of water.3. Water insoluble solids and liquid should be placed in the waste container provided.4. Chromium ion in the +6 oxidation state should be reduced to the +3 state with a mild

reducing agent before disposal.

8/19/2019 Lab_Manual_STK1211_2015_2016

7/38

7

Apparatus and techniques

The following is a summary of the basic analytical laboratory techniques and equipment youwill use for this semester. Proper techniques are essential as acceptable error in aquantitative chemical analysis is seldom greater than 0.1 % . There are a number of “hard andfast” rules presented that must be followed to minimize any hazards to yourself, your lab -coworkers, and the lab equipment. Read this section at the beginning of the semesterand refer to any of this material as often as necessary.

I. “Clean” and “Clean and Dry” glassware

You will notice throughout the semester that you are asked to use “Clean” glassware at timesand “Clean and Dry” glassware at other times. “Clean” glassware may be wet with yoursolvent (usually distilled water), and many times it is not worth the effort to dry the piece ofglassware.

In general, if you want to maintain the concentration of the solution being transferred, youwill want the final container to be “Clean and Dry”. However, if you are only c oncerned about

the amount of the compound being transferred, the final container need only be “Clean”.

“Clean” glassware means that all compounds and materials have been washed out. The finalwashing should be with your solvent. In analytical chemistry, the solvent is defined as theliquid or solution that you would use to dilute the solution in question, usually distilled water.

II. Desiccators and Handling Dried Compounds

When using primary standards (compounds that are presumed to be 100% pure) in analyses,it is essential that there are no crystal waters present so that the mass of the primarystandard measured on the balance is equal to the actual mass. Typically, compounds are driedby placing them in an oven at 105 to 120 C. After several hours, all (or most) crystal watershave been driven off. However hot compounds cannot be accurately weighed ( all itemsweighed on a balance must be at room temperature ), so there must be a way to cool adried compound without re-exposing it to water vapor in the atmosphere.

Desiccators are containers designed to prevent the re-hydration of solids. The bottom half ofthe desiccator is filled with an anhydrous salt, such as calcium chloride. The dried compoundand its container sit in the top half, which is separated from the bottom half by a grid orscreen. The desiccator lid can be sealed with vacuum grease to prevent water vapor fromseeping inside.

Always cool a dried compound in a desiccator before weighing. A dried compound can be keptin the desiccator if that compound has to be available throughout the lab period.

III. Electronic and Analytical BalancesElectronic balances are quite simple to use. As you probably know, the "tare" button resets themass reading at zero, and there is usually another button (sometimes labeled "cal" forcalibrate) to set the mass units. You will always want the mass units in grams. Becauseelectronic balances are fragile, you need to observe the following guidelines.

1. Always clean the balance after using it -- use a soft brush or Kimwipe to remove anyextraneous material from the balance pan.

2. All items and compounds placed on the balance must be at room temperature -- this cancome into play when weighing dried compounds. Cool dried compounds in a desiccatorbefore weighing them on the balance.

3. If you are making repeated weighing in the same container, it is recommended that youalways tare the empty balance and record the mass of the empty container. Then, recordthe mass of the container with sample, and calculate the mass of the sample by difference.

8/19/2019 Lab_Manual_STK1211_2015_2016

8/38

8

4. If you are instructed to "accurately weigh" something, use a balance with 4 decimal places.This is referred to as analytical balance. The maximum capacity of an analytical balanceis usually small (60 or 180 g), therefore use only weighing boats or weighing paper ascontainers on the balance. On the electronic balance with 3 decimal places, you can oftenuse small beakers or flasks as containers.

IV. Volumetric Flasks and Quantitative Transfers

Volumetric flasks are calibrated to contain an exact volume of solution when the solution levelis exactly at the mark on the neck of the flask (the bottom of the meniscus should lie exactlyat this mark). Note the following rules in handling volumetric flasks.

1. To clean volumetric flasks -- Each washing should have a volume that is about 10 to20% of the capacity of the vol. flask. Typically, you should wash the flask 3 times withdilute acid (e.g., 1 to 6 M HCl), 3 times with distilled water, and 3 times with your solvent(if it is not distilled water). You can skip the acid washings if you have no solid residue inthe flask.

2. Never heat a volumetric flask -- heating causes the glass to expand, changing thevolume it contains.

3. NEVER place a solid directly into a volumetric flask -- What do you do if you fill theflask to the mark and the solid won't dissolve? Dissolve the solid in another container andquantitatively transfer the solution to the volumetric flask. For example, if you want tomake a 100-mL solution of NaCl, dissolve the NaCl in a beaker with 75 to 90 mL of waterand then transfer this to the volumetric flask.

4. Never shake a volumetric flask -- when mixing a solution in a vol. flask, gently invertthe flask 8 to 10 times.

5. Quantitative transfers are vital to accurate analyses. In simple terms, quantitatively

transferring something means washing the original container and all glassware involvedin the transfer with solvent and adding those washings to the final container, usually avolumetric flask. Here are some guidelines to transferring solutions and solids tovolumetric flask.

Pour the solution through a funnel into the volumetric flask. Wash the original container witha small amount of solvent and pour the washing through the funnel into the volumetric flask.Repeat the washing 2 or 3 times if possible. Dilute the solution in the volumetric flask to themark.

V. Volumetric Pipettes

Volumetric pipettes are calibrated to deliver an exact volume of liquid or solution. Volumetric

pipettes have only one calibration mark. You may have seen graduated pipettes that havecalibration marks throughout the length of the pipette, but these are far less accuratethan volumetric pipettes. To fill a pipette, simply draw in liquid to the mark. Usually it iseasiest to initially overshoot the mark and then let the liquid drain from the pipette until thebottom of the meniscus lies exactly on the calibration mark.

8/19/2019 Lab_Manual_STK1211_2015_2016

9/38

9

Note the following rules in handling pipettes.

1. Never pipette with your mouth! Always use a rubber bulb, regardless of what you aretaught in biology classes.

2. Always clean a pipette before its initial use -- For each washing, draw liquid into thepipette so that the bulb is 1/4 to 1/2 full (this can be less for large pipettes). Carefully swirlthe liquid throughout the inside of the pipette (don't let liquid pour out the top!), and letthe liquid drain from the pipette. Wash the pipette 3 times with distilled water and 3times with the solution you are going to transfer (not just the solvent). If you are using the

same pipette for different solutions, you need to repeat the washing procedure every timeyou switch solutions.

3. Never force liquid out of a pipette -- always let the liquid drain of its own accord. Thecalibration mark takes into account any liquid that is retained in the tip of the pipette.When the liquid has stopped draining, touch the tip of the pipette against the side of thecontainer to release any hanging drops.

4. At the end of a lab period, always wash used pipettes 3 times with distilled water.

VI. Burettes

Burettes allow you to accurately deliver volumes of liquid that cannot be measured byvolumetric pipettes or micropipettors. The proper use of burettes is essential to accuratetitration analyses. In several cases, an analyte can be determined more accurately using atitrimetric method rather than an instrumental method -- it's just that the convenience of theinstrumental use is sometimes the deciding factor in choosing an analytical method. Note thefollowing rules and guidelines in using burettes.

1. To clean a burette -- fill the burette with distilled water and drain a large portion of it tosee if any water adheres to the inside walls of the burette. If so, clean the burette with afew milliliters of soap solution and a burette brush, and wash the burette with three 5-mLportions of tap water. When washing the inside of a burette, pour about 5 mL of liquidinto the burette with the stopcock closed. Carefully swirl the liquid for a few seconds sothat it comes in contact with the entire inside surface area of the burette, and pour theliquid out the top. If no water adheres to the inside walls of the burette, proceed to wash it

3 times with distilled water and 3 times with the solution with which you are going totitrate your sample.

8/19/2019 Lab_Manual_STK1211_2015_2016

10/38

10

2. Getting rid of air bubbles -- fill the clean burette with your solution. There will be airbubbles inside the stopcock, and you must remove these (you can't measure the volume ofan air bubble in a burette, so if an air bubble pops out in the middle of your titration,you're sunk). It's usually easiest to force bubbles out the stopcock, so simply open thestopcock and let the solution drain until the air bubbles are removed. If this isn't working,you can try gently tapping the base of the burette while the solution is draining. Thismay force the air bubbles to rise through the solution.

3. Always record volume levels to the nearest 0.01 mL -- Although the calibration marksare only at every 0.1 mL, you can always estimate the extra decimal place.

For Example:

The reading in this picture is between 8.2 and 8.3 mL. The second decimal place would beestimated to be about 0.07 mL, giving a reading of 8.27 mL .

To obtain readings, it helps to hold a white card behind the burette. Note reading at aparticular part of meniscus and always measure at this part.

4. NEVER record an initial volume level of 0.00 mL -- since there are no calibrationmarks above 0.00 mL, you have no upper reference with which to base any possible errorin your reading.

5. Using your wash bottle, you can add "half-drops" to your titration flask (see above figure).Run a drop of solution part-way out of your burette tip. Squirt distilled water on this half-drop and into your titration flask. This procedure is perfectly legitimate because youaren't worried about the concentration of anything in your titration flask.

6. At the end of an experiment, always wash your burette 3 times with distilled water .

8/19/2019 Lab_Manual_STK1211_2015_2016

11/38

8/19/2019 Lab_Manual_STK1211_2015_2016

12/38

12

REAGENTS

Sodium hydroxide pelletsPotassium hydrogen phthalate (KHP)Phenolphthalein

Unknown vinegar

PROCEDURE

Part A: Preparation of the Sodium Hydroxide Solution

1. Clean and rinse a 1-L volumetric flask and stopper. Label the flask “Approx. 0.1 M NaOH”. Put about 500 mL of distilled water into the flask.

2. Weigh out approximately 4.0 g of sodium hydroxide pellets and transfer to the 1-Lflask. Stopper and shake the flask to dissolve the sodium hydroxide.

3. When all the sodium hydroxide pellets have dissolved, add additional distilled waterto the bottle until the mark on the neck of the flask. Stopper and shake thoroughly tomix.

Part B: Standardization of the Sodium Hydroxide Solution

1. Set up the burette in the burette clamp. Rinse and fill the burette with the sodiumhydroxide solution just prepared.

2. Clean three 250-mL Erlenmeyer flasks with water, and then rinse with distilled water.Label them as 1, 2, and 3

3. Remove the bottle of dried KHP from the oven. When the KHP is completely cool,

weigh three samples of KHP between 0.6 and 0.8 g , one for each of the Erlenmeyerflasks. Record the exact weight of each KHP sample to the nearest mg ( 0.001 g).4. Add 100 mL of distilled water to KHP sample 1. Add 2 – 3 drops of phenolphthalein

indicator solution. Swirl to dissolve the KHP sample completely.5. Record the initial reading of the NaOH solution in the burette to the nearest 0.02 mL.6. Add NaOH solution from the burette to the sample in the Erlenmeyer flask, swirling

the flask constantly during the addition.7. When the titration approaching the endpoint, add NaOH one drop at a time, with

constant swirling, until one single drop of NaOH causes a permanent pale pink colourthat does not fade on swirling. Record the reading of the burette to the nearest 0.02mL.

8. Repeat step 4 – 7 with the other 2 KHP samples.9. Given that the molecular mass of KHP is 204.2, calculate the number of moles of KHP

in samples 1, 2, and 3.10. From the number of moles of KHP present in each sample, and from the volume of

NaOH solution used to titrate the sample, calculate the molar concentration (M) of NaOH in the titrant solution. The reaction between NaOH and KHP is of 1 : 1stoichiometry.

8/19/2019 Lab_Manual_STK1211_2015_2016

13/38

13

Part C: Analysis of a Vinegar Solution

Vinegar is a dilute solution of acetic acid and can be effectively titrated with NaOH usingthe phenolphthalein endpoint.

1. Clean three Erlenmeyer flasks, and label as samples 1, 2, and 3.2. Rinse the 5-mL pipette with small portions of the vinegar solution and discard the

rinsings.3. Using the pipetter, pipette 5 mL of the vinegar solution into each of the Erlenmeyer

flasks. Add about 100 mL of distilled water and 2 – 3 drops of phenolphthaleinindicator solution to each flask.

4. Refill the burette with the NaOH solution and record the initial reading of the buretteto the nearest 0.02 mL. Titrate Sample 1 of vinegar in the same manner as in thestandardization until one drop of NaOH causes the appearance of the pink colour.

5. Record the final reading of the burette to the nearest 0.02 mL.6. Repeat the titration for the other two vinegar samples.7. Based on the volume of vinegar sample taken, and on the volume and average

concentration of NaOH titrant used, calculate the molar concentration of the vinegarsolution.

8. Given that the formula mass of acetic acid is 60.0, and the density of the vinegarsolution is 1.01 g/mL, calculate the percent by mass of acetic acid in the vinegarsolution.

QUESTIONS

1. Give the definition of indicators.

2. Suppose a NaOH solution were to be standardized against pure solid primary standardgrade KHP. If 0.4538 g of KHP requires 44.12 mL of the NaOH to reach a

phenolphthalein endpoint, what is the molarity of the NaOH solution?

3. Commercial vinegar is generally 5.0 0.5% acetic acid by weight. Assuming this to bethe true value for your sample, by how much were you in error in your analysis?

8/19/2019 Lab_Manual_STK1211_2015_2016

14/38

14

Experiment 1: Acid and Base Titrations

Data Sheet

Name:

Student No.:

Date: _______________________________________________________________________

Part A: Standardization of the Sodium Hydroxide Solution

Particulars Trial 1 Trial 2 Trial 3

Mass of KHP taken (g)

Final burette reading (mL)

Initial burette reading (mL)

Volume of NaOH used (mL)

Molarity of NaOH solution

Average molarity of NaOH solution

Part B: Analysis of a Vinegar Solution


Volume of vinegar solution used (mL)



Volume of NaOH used (mL)

Molarity of NaOH solution

Molarity of vinegar solution

% mass of acetic acid in vinegar

Average molarity of vinegar solution

Average % mass of acetic acid in vinegar

8/19/2019 Lab_Manual_STK1211_2015_2016

15/38

15

EXPERIMENT 2: WATER ANALYSIS: SUSPENDED SOLIDS ANDDISSOLVED OXYGEN

INTRODUCTION

Water is a very important resource for human being. A large portion of our body is water.A healthy human being can live for weeks without foods, but will not survive for dayswithout water. When you are thirsty and need a drink of water, do you take for grantedthat the tap water is safe to drink? Do you think about what might be dissolved orsuspended in the water?

Water in the environment has a large number of impurities. Dissolved solids are water-soluble substances, usually salts. Naturally occurring dissolved solids generally resultfrom the movement of water over mineral deposits, such as limestone. These dissolvedsolids generally consist of the sodium, calcium, magnesium, and potassium cations and the

chloride, sulfate, bicarbonate, carbonate, bromide, and fluoride anions. The dissolvedsolids are responsible for the “hard” water that exists in some locales. Anthropoge nic(human-related) dissolved solids include nitrates from fertilizer runoff and human wastes,

phosphates from detergents and fertilizers, and organic compounds from pesticides andindustrial wastes. Salinity, a measure of the dissolved solids in a water sample, is definedas the grams of dissolved solids per kilogram of water.

Suspended solids are very finely divided particles that are water insoluble but arefilterable. These particles are kept in suspension by the turbulent action of the movingwater. Examples of suspended solids include decayed organic matter, sand, salt, and clay.Total solids are the sum of the dissolved and suspended solids in the water sample. In thisexperiment the total solids and the dissolved solids are determined directly; the suspendedsolids are assumed to be the difference, since total solids = di ssolved soli ds + suspendedsolids

The dissolved oxygen is determined based on the iodine/thiosulfate method. Excess potassium iodide is added to the acidified water sample, the dissolved oxygen (or otheroxidizing agents that are present) will oxidize the iodide ions to iodine (in the form oftriiodide, I 3 ). The iodine formed will then be determined by titration with the standardsolution of sodium thiosulfate with starch solution as the indicator.

APPARATUS

Burette with standPipetteBeaker250 mL Erlenmeyer flasksEvaporation dishesWatch glassFilter funnelFilter paper

REAGENTS

8/19/2019 Lab_Manual_STK1211_2015_2016

16/38

16

KI0.05M Na 2S2O3 solutionStarch solution6M Sulfuric acid

PROCEDURE:

I: Determination of Suspended Solids in a Water Sample

1. Clean, dry, and determine the mass (to the nearest milligram, 0.001 g) of twoevaporation dishes. Be certain that you can identify each. Use the same balance forthe remainder of the experiment.

Part IA: Determination of dissolved solids

1. Gravity filter about 50 mL of a thoroughly stirred or shaken water sample into aclean, dry 100 mL beaker. While waiting for the filtration to be completed, proceed toPart B.

2. Pipette a 25 mL portion of the filtrate into one of the evaporating dishes. Determinethe mass of the combined evaporating dish and sample. Place the dish containing thesample on the wire gauze and heat slowly (do not boil) the mixture to dryness. As themixture nears dryness, cover with a watch glass, and reduce the intensity of theflame. If spattering does occur, allow the dish to cool to room temperature, rinse theadhered solids from the watch glass and return the rinse to the dish.

3. Again heat slowly, and avoid further spattering. After all of the water has evaporated,maintain a small flame beneath the dish for 3 minutes. Allow the dish to cool to roomtemperature and determine its final mass.

Part IB: Determination of total solids and suspended solids

1. Thoroughly agitate 100 mL of sample; pipette a 25 mL aliquot of this sample into thesecond evaporating dish. Evaporate the sample to dryness as described in Part A.

2. Calculate the mass of total and suspended solids in the original water sample, usingdata from Part A.

Part IC: Analysis of data

1. Compare your data with three other groups in your laboratory who analyzed the samewater sample. Record their results on the data sheet.

2. Calculate the standard deviation of the suspended solids from the four analyses on thewater sample.

II: Determination of Dissolved Oxygen in a Water Sample

1. Prepare three clean Erlenmeyer flasks.

2. Clean a burette with tap water, then rinse with distilled water, and finally rinse withthe 0.05M sodium thiosulfate standard solution. Then fill the burette with the sodiumthiosulfate solution and record its initial volume to the nearest 0.02 mL

8/19/2019 Lab_Manual_STK1211_2015_2016

17/38

17

3. Using a clean pipette, transfer 25 mL of the water sample into each of the Erlenmeyerflasks and label as sample 1, 2 and 3.

4. Add about 2 g of KI and 10 mL of 6M sulfuric acid to sample 1. Iodide ions will beoxidized to iodine, and then it is titrated with the standard sodium thiosulfatesolution.

5. On titration, the colour of sample will become lighter. Stop adding the Na 2S2O3 solution when the colour of the sample becomes light yellow that means it has nearlyreached the end point. Now add 2 – 3 mL of starch solution to the sample. Thesample will appear in dark-blue colour.

6. Continue adding the Na 2S2O3 solution drop by drop until the dark-blue colour justdisappears. Record the final volume to the nearest 0.02 mL

7. Repeat the titration with the sample 2 and 3. From the concentration and the volumeof the Na 2S2O3 solution used in the titration, calculate the concentration of oxygen

present in the water sample.

QUESTIONS

1. In evaporating a solution to dryness in an evaporating dish, why must the heating rate be decreased as the mixture nears dryness?

2. How do pollutants such as sewage lower the dissolved oxygen content of watersources?

3. A 25 mL aliquot of a well-shaken sample of river water is pipetted into 27.211 g

evaporating dish. After the mixture is heated to dryness, the dish and remainingsample has a mass of 43.617 g. Determine the total solids in the sample, express inunits of g/kg sample. Assume the density of the sample to be 1.01 g/mL.

8/19/2019 Lab_Manual_STK1211_2015_2016

18/38

18

Experiment 2: Water Analysis: Suspended Solids and Dissolved Oxygen

Data Sheet

Name:

Student No.:

Date:

______________________________________________________________________________________

I: Determination of Suspended Solids in a Water Sample

Part IA: Determination of dissolved solids

Particular DataMass of evaporating dish (g)

Mass of evaporating dish + water sample (g)

Mass of water sample (g)

Mass of evaporating dish + dried sample (g)

Mass of dissolved solids in 25 mL aliquot (g)

Mass of dissolved solids per total mass of sample (gsolids/g sample)

Dissolved solids (g solids/kg sample)

Part IB: Determination of total solids and suspended solids

Particular Data

Mass of evaporating dish (g)

Mass of evaporating dish + water sample (g)

Mass of water sample (g)

Mass of evaporating dish + dried sample (g)Mass of total solids in 25 mL aliquot (g)

Mass of total solids per total mass of sample (gsolids/g sample)

Suspended solids (g solids/kg sample)

8/19/2019 Lab_Manual_STK1211_2015_2016

19/38

19

Part IC: Analysis of data

Particular Group 1 Group 2 Group 3 Group 4

Dissolved solids (g/kg)

Total solids (g/kg)Suspended solids (g/kg)

Average Suspended solids (g/kg)

II: Determination of Dissolved Oxygen in a Water Sample

Concentration of Na 2S2O3 standard solution: _______________________________

Particular Trial 1 Trial 2 Trial 3



Volume of S 2O32 used (mL)

Moles of S 2O32 used (mol)

Volume of water sample (mL)

Concentration of sample (M)

Mean concentration of sample (M)

8/19/2019 Lab_Manual_STK1211_2015_2016

20/38

20

EXPERIMENT 3: WATER ANALYSIS: HARDNESS OF WATERAND CHLORIDE ION/CHLORINE IONCONCENTRATION

INTRODUCTION

What does it mean when we say that water is “hard”? Hard water contains the dissolvedsalts of calcium, magnesium, and iron ions which are called hardening ions. In lowconcentrations these ions are not considered harmful for domestic use, but at higherconcentrations these ions interfere with the cleansing action of soaps by forming insolublecompounds with soaps. Soaps, sodium salts of fatty acids such as sodium stearate,C17H35CO 2 Na, are very effective cleansing agents so long as they remain soluble; the

presence of the hardening ions however causes the formation of a grey, insoluble soapscum such as (C 17H35CO 2)2Ca. This grey precipitate appears as a “bathtub ring” and it alsoclings to clothes, causing white clothes to appear grey.

In industries, hard water can accelerate the corrosion of steel pipes, especially thosecarrying hot water. It is also responsible for the formation of “boiler scale” on tea kettlesand pots used for heating water. The boiler scale is a poor conductor of heat and thusreduces the efficiency of transferring heat. Boiler scale also builds on the inside of hotwater pipes to decrease the flow of water; in extreme cases, this buildup causes the pipe to

break. Boiler scale consists primarily of the carbonate salts of the hardening ions. Groundwater becomes hard as it flows through underground limestone (CaCO 3) deposits; Surfacewater similarly accumulates hardening ions as a result of it flowing over limestonedeposits. Because of the relative large natural abundance of limestone deposits and othercalcium minerals, it is not surprising that Ca 2+ ion, in conjunction with Mg 2+ , is a major

component of the dissolved solids in water.

Hardness Classification of Water

Hardness (ppm CaCO 3) Classification

< 15 ppm Very soft water

15 ppm – 50 ppm Soft water

50 ppm – 100 ppm Medium hard water

100 ppm – 200ppm Hard water

> 200 ppm Very hard water

The concentration of the hardening ions in a water sample is commonly expressed asthough the hardness is due exclusively to CaCO 3. The unit for hardness is mg CaCO 3/L,which is also ppm CaCO 3.

In this experiment a titration technique is used to measure the combined Ca 2+ and Mg 2+ concentration in a water sample. The titrant is EDTA, the disodium salt ofethylenediaminetetraacetic acid (abbreviated Na 2H2Y).

In aqueous solution Na 2H2Y dissociates into Na+ and H 2Y

2 ions. The H 2Y2 ion reacts

with the hardening ions, Ca 2+ and Mg 2+ , to form very stable complex ions, especially in a

solution buffered at a pH of about 10. An ammonia-ammonium ion buffer is often used forthis pH adjustment in the analysis. A special indicator called Eriochrome Black T (EBT) is

8/19/2019 Lab_Manual_STK1211_2015_2016

21/38

21

used to detect the endpoint in the titration. EBT forms complex ions with Ca 2+ andMg 2+ions, but binds more strongly to Mg 2+ ions. Because only a small amount of EBT isadded, only a small quantity of Mg 2+ is complexed; no Ca 2+ ion is complexed to EBT,therefore most of the hardening ions remain “free” in solution. The EBT indicator is bluein solution but the [Mg-EBT] 2+ complex ion is red.

Mg 2+(aq ) + EBT( aq ) [Mg-EBT] 2+(aq ) blue red

There even before any H 2Y2 titrant is added for the analysis, the solution is red. As H 2Y

2 titrant is added, it complexes with the “free” Ca 2+ and Mg 2+.

Ca 2+(aq ) + H 2Y2 (aq ) CaY 2 (aq ) + 2H +(aq )

Mg 2+(aq ) + H 2Y2 (aq ) MgY 2 (aq ) + 2H +(aq )

Once the H 2Y2 complexes all of the “free” Ca 2+ and Mg 2+ from the water sample, it thenremoves the Mg 2+ from the [Mg-EBT] 2+ complex; the solution turns from red back to bluecolour of the indicator, and the endpoint is reached.

[Mg-EBT] 2+(aq ) + H 2Y2 (aq ) MgY 2 (aq ) + 2H +(aq ) + EBT( aq )

red blue

For the endpoint to appear Mg 2+ must be present; therefore a small amount of MgY 2 isusually added to the buffer solution. The added Mg 2+ does not affect the amount of H 2Y

2 used in the analysis because an equimolar amount of Na 2H2Y is also added.

The concentration of chloride ions in water sample can be determined by Mohr precipitation titration. The water sample is titrated with standard silver nitrate solution andsodium chromate as indicator. A small amount of calcium carbonate is added to changethe pH of the water sample to basic condition.

APPARATUS

Volumetric flaskBurette with standPipetteBeaker250 mL Erlenmeyer flasks

REAGENTS

Na 2EDTAStandard Ca 2+ solutionMg/EDTA solutionAmmonia-ammonium ion buffer solutionEriochrome Black T0.05M AgNO 3 solutionSodium chromate, Na 2CrO 4 Calcium carbonate, CaCO 3

8/19/2019 Lab_Manual_STK1211_2015_2016

22/38

22

PROCEDURE

I: Determination of the hardness of a water sample

Part IA: To prepare a standard 0.01M Na 2EDTA solution

1. Measure about 1.25 g ( 0.01 g) of Na 2EDTA (molar mass = 372.24 g/mol); transfer itto a 250 mL volumetric flask containing about 200 mL of distilled water and stir todissolve. Dilute to the “mark” on the volumetric flask with distilled water.

2. Prepare a burette for titration. Rinse a burette with the Na 2EDTA solution and then fill.Record the initial volume ( 0.02 mL) of the solution.

3. Pipette out 25.0 mL of the standard Ca 2+ solution provided into a 250 mL Erlenmeyerflask, and record its molar concentration. Then add 2 mL of the buffer (pH = 10)solution, 5 mL of Mg/EDTA solution, and 5 – 6 drops of EBT indicator. Titrate thestandard Ca 2+ solution with the Na 2EDTA solution; swirl continuously. Near theendpoint, slow the rate of addition to drops; the last few drops should be added at 3 – 5seconds intervals. The solution changes from red to purple to blue; the solution is blueat the endpoint.

4. Repeat the titration twice, then calculate the concentration of the Na 2EDTA solution.

Part IB: Analysis the hardness of water sample

1. Obtain about 100 mL of a water sample. If the water sample is too turbid, you willneed to gravity filter the sample before the analysis, and if the sample is acidic, add1M NH 3 until it is basic to litmus.

2. Pipette out 25 mL of the water sample into a 250 mL Erlenmeyer flask, add 2 mL ofthe buffer (pH = 10) solution, 5 mL of Mg/EDTA solution, and 5 – 6 drops of EBTindicator, then titrate with the Na 2EDTA solution till the endpoint is reached. Repeat

the titration thrice and then determine the hardness of the water sample.

II: Determination of the concentration of chloride ion

1. Rinse a clean burette with some standard 0.05M AgNO 3 solution, and then fill. Recordthe initial volume ( 0.02 mL) of the solution.

2. Pipette 25 mL of the water sample into an Erlenmeyer flask. Add 3 – 5 drops ofsodium chromate solution (yellow colour) and a piece of pea size calcium carbonate.

3. Titrate the water sample slowly with the standard 0.05M AgNO 3 solution from the burette with constant swirling. At the beginning, a white precipitate, AgCl is formed.As soon as all the Cl ions have been precipitated, a red precipitate, Ag 2CrO 4, starts

being formed. Stop the titration when the endpoint is reached, that is the red colourremains.

4. Repeat the titration twice and then determine the concentration of the chloride ions present in the water sample.

QUESTIONS

1. Why was it necessary to add a small amount of magnesium/EDTA complex to thecalcium sample before titration?

2. A 25 mL sample of well water for chloride determination requires 34.32 mL of0.05012 M AgNO 3 solution to reach a sodium chromate endpoint. Calculate theconcentration of chloride ion in the water sample.

8/19/2019 Lab_Manual_STK1211_2015_2016

23/38

23

Experiment 3: Water Analysis: Hardness Of Water And ChlorideIon/Chlorine Ion Concentration

Data Sheet

Name:

Student No.:

Date: _______________________________________________________________________

I: Determination of the hardness of a water sample

Part IA: To prepare a standard 0.01M Na 2EDTA solution

Molarity of Ca 2+ solution: ________________________________


Volume of standard Ca 2+ solution used (mL)



Volume of Na 2EDTA used (mL)

Molarity of Na 2EDTA solution (M)

Average molarity of Na 2EDTA solution (M)

Part IB: Analysis the hardness of water sample




Volume of Na 2EDTA used (mL)

Moles of Na 2EDTA used (mol)




8/19/2019 Lab_Manual_STK1211_2015_2016

24/38

24

II: Determination of the concentration of chloride ion/chlorine ion


Final burette reading (mL)Initial burette reading (mL)

Volume of AgNO 3 used (mL)

Moles of AgNO 3 used (mol)




8/19/2019 Lab_Manual_STK1211_2015_2016

25/38

25

EXPERIMENT 4: REDOX TITRATION – ASCORBIC ACID

INTRODUCTION

The human body does not synthesize vitamins; therefore the vitamins that we need forcatalyzing specific biochemical reactions are gained only from the food that we eat. Weare generally aware that vitamin C can be obtained from citrus fruits, but it can also beobtained from a variety of fresh fruits and vegetables. However, storage and processingcauses vegetables to lose a part of their vitamin C content. Cooking leaches the water-soluble vitamin C from the vegetables; in addition, the high temperatures accelerate itsdegradation by air oxidation. Therefore to maximize the intake of vitamin C, only freshfruits or vegetables should be consumed.

Vitamin C, also called ascorbic acid, is one of the more abundant and easily obtainedvitamins in nature. It is a colourless, water-soluble acid that, in addition to its acidic

properties, is a powerful biochemical reducing agent, meaning it readily undergoesoxidation, even from the oxygen of the air.

Even though ascorbic acid is an acid, its reducing properties are used in this experiment toanalyze its concentration in various samples. There are many other acids present in foodsthat would interfere with an acid analysis and not permit us to selectively determine theascorbic acid content. The equation for the oxidation of ascorbic acid is

C6H8O6(aq) + H 2O(l) C6H8O7(aq) + 2H +(aq) + 2e

In analysing for ascorbic acid, the sample is dissolved in water and treated with ameasured excess of iodate ion, IO 3 , a strong oxidizing agent; in an acidic solutioncontaining an excess of iodide ion, I , IO 3 converts to I 3 (red-brown), a milder oxidizingagent (Equation 1).

IO3 (aq) + 8I (aq) + 6H +(aq) 3I3 (aq) + 3H 2O(l) ------ (Equation 1)

Some of the I 3 then oxidizes the ascorbic acid (Equation 2) present in the sample.

I3 (aq) + C 6H8O6(aq) + H 2O(l) C6H8O7(aq) + 3I (aq) + 2H +(aq) ---- (Equation 2)

The remaining I 3 (the excess, “xs”) is titrated with a standard thiosulfate, S 2O32 , solution,

producing the colourless I and S 4O62 ions (Equation 3).

(xs)I 3 (aq) + 2S 2O32 (aq) 3I (aq) + S 4O6

2 (aq) ------ (Equation 3)

Therefore the difference between the I 3 generated initially and that which is titrated inexcess is a measure of the ascorbic acid content of the sample. The stoichiometric point isdetected using starch as an indicator. Just prior to the disappearance of the red-brown I 3 inthe titration, starch solution is added; this forms the deep-blue ion, [I 3·starch] . Theaddition of the S 2O3

2 titrant is continued until the [I 3·starch] has been reduced to I , thesolution appears colourless at the end point.

8/19/2019 Lab_Manual_STK1211_2015_2016

26/38

26

In this experiment you are required to prepare and standardize a Na 2S2O3 solution usingsolid KIO 3 as a primary standard. The standard solution is then used to analyze forascorbic acid in the samples provided.

APPARATUS

250-mL volumetric flask250-mL Erlenmeyer flask50-mL burette with standGlass stirring rodBeakerMuslinFilter funnel

REAGENTSPotassium iodate, KIO 3 Sodium thiosulfate, Na 2S2O3 Potassium iodide, KIStarch solution0.5 M H 2SO 4 Sodium hydrogen carbonate, NaHCO 3 Vitamin C tabletsFresh fruit juices

PROCEDURE

Part A: A Primary Standard 0.01 M Potassium Iodate, KIO 3, Solution

1. Measure about 0.5 g ( 0.001 g) of KIO 3 on weighing paper, transfer the solid to a 250mL volumetric flask, dissolve and dilute to the mark.

2. Calculate and record the molar concentration of the KIO 3 solution.

Part B: A Standard 0.1 M Na 2S2O 3 Solution

1. Dissolve about 6 g ( 0.001 g) of Na 2S2O3·5H 2O with distilled water and dilute to 250

mL. Stir until the salt dissolves.2. Properly prepare a clean 50 mL burette and fill it with the Na 2S2O3 solution, drain the

tip of air bubbles, and read and record the initial volume ( 0.02 mL).3. Pipette 25 mL of the standard KIO 3 solution into a 250 mL Erlenmeyer flask and add

about 1 g ( 0.01 g) of solid KI. Add about 5 mL of 0.5 M H 2SO 4 and 0.1 g of NaHCO 3 (The NaHCO 3 reacts in the acidic solution to produce CO 2 , providing an inertatmosphere above the solution and minimizing the possibility of the air oxidation of Iions).

4. Immediately begin titrating with the Na 2S2O3 solution. When the red-brown solution(due to I 3 ) changes to a pale yellow colour, add 2 mL of starch solution. Stirring

constantly, continue titrating slowly until the blue colour disappears.

8/19/2019 Lab_Manual_STK1211_2015_2016

27/38

27

5. Repeat the procedure twice by rapidly adding the Na 2S2O3 solution until 1 mL beforethe endpoint point. Add the starch solution and continue titrating until the solution iscolourless.

6. Calculate and record the molar concentration of the Na 2S2O3 solution.

Part C: Sample Preparation

(a) Vitamin C tablet

1. Read the label to determine the approximate mass of vitamin C in each tablet. Measure( 0.001 g) the fraction of the total mass of a tablet that corresponds to 100 mg ofascorbic acid.

2. Dissolve it in a 250 mL Erlenmeyer flask with 40 mL of 0.5 M H 2SO 4 (Remember thatvitamin tablets contain binders and other material that may be insoluble in water – donot heat in an attempt to dissolve it) , and then add about 0.5 g NaHCO 3.

3. Proceed immediately to Part D.

(b) Fresh fruit sample

1. Filter about 120 mL of freshly squeezed juice through several layers of muslin.2. Measure the mass ( 0.01 g) of a clean, dry 250 mL Erlenmeyer flask. Add about 100

mL of the filtered juice and again determine the mass.3. Add 40 mL of 0.5 M H 2SO 4 and 0.5 g NaHCO 3, and then proceed immediately to Part

D.

Part D: Vitamin C analysis

1.

Pipette 25.0 mL of the standard KIO 3 solution (from Part A) into the sample solutionfrom Part C and add 1 g of KI.

2. Add about 5mL of 0.5 M H 2SO 4 and 0.1 g NaHCO 3. Titrate the excess I 3 in thesample with the standard Na 2S2O3 solution as described in Part B.4. Read and recordthe final burette reading ( 0.02 mL).

3. Repeat the analysis twice in order to complete the three trials.4. Calculate the percent of ascorbic acid in the sample.

QUESTIONS

1. Explain why cooked fruits and vegetables have lower vitamin C content than freshfruits and vegetables.

2. If the blue colour does not appear when the starch solution is added during thetitration, should you continue titrating or discard the sample? Explain.

3. 177.42 mL of a fruit juice contains 32% of the recommended daily allowance ofvitamin C (equal to 60 mg). How many mL of the fruit juice will provide 100% of therecommended daily allowance?

8/19/2019 Lab_Manual_STK1211_2015_2016

28/38

28

Experiment 4: Redox Titration – Ascorbic Acid

Data Sheet

Name:

Student No.:

Date: _______________________________________________________________________

Part A: A Primary Standard 0.01 M Potassium Iodate, KIO 3, Solution

Mass of KIO 3 (g)

Moles of KIO 3 (mol)

Molarity of standard KIO 3 solution (M)

Part B: A Standard 0.1 M Na 2S2O 3 Solution


Volume of KIO 3 solution (mL)

Moles of KIO 3 titrated (mol)

Moles of I 3 generated (mol)



Volume of Na 2S2O3 added (mL)

Moles of Na 2S2O3 added (mol)

Molarity of Na 2S2O3 solution (M)

Average molarity of Na 2S2O3 solution (M)

8/19/2019 Lab_Manual_STK1211_2015_2016

29/38

29

Part C & D: Sample Preparation & Vitamin C analysis

Sample name: ________________________________________


Mass of sample (g)Volume of KIO 3 solution added (mL)

Moles of KIO 3 added (mol)

Total moles of I 3 generated (mol)



Volume of Na 2S2O3 added (mL)

Moles of S 2O32

added (mol)Moles of I 3 titrated with S 2O3

2 (mol)

Moles of I 3 reduced by C 6H8O6 (mol)

Moles of C 6H8O6 in sample (mol)

Mass of C 6H8O6 in sample (g)

Percent of C 6H8O6 in sample (%)

Average percent of C 6H8O6 in sample (%)

8/19/2019 Lab_Manual_STK1211_2015_2016

30/38

30

EXPERIMENT 5: GRAVIMETRIC DETERMINATION OFNICKEL(II) ION

INTRODUCTION

The percentage of nickel in a sample may be determined gravimetrically by precipitationwith the organic reagent dimethylglyoxime (DMG). DMG is a bidentate ligand. Nickel isoften added in small amount during the production of steel; the precipitation analysis withDMG is especially useful in this situation. DMG is a complexing agent, forming acharacteristic bright red coordination compound with nickel ion. In addition to thegravimetric determination of nickel, DMG is also often used as a spot test to detect the

presence of nickel in a sample on a qualitative basis. DMG does precipitate a few othermetal ions like platinum, but the bright red colour of the Ni(II)/DMG precipitate is distinctfrom others.The precipitate produced by DMG with nickel(II) ion contains two DMG speciescomplexed per nickel ion and consists of 20.32% nickel by weight. The precipitate is veryfluffy and has a very low density, which makes it somewhat difficult to handle in any greatquantity. For this reason, the practical use of DMG in nickel analyses is restricted tosamples in which the percent of nickel is rather small. As an organic reagent, DMG is notvery soluble in water and is provided as a solution in alcohol. The presence of volatilealcohol, combined with the nature of the Ni/DMG precipitate itself, causes the precipitateto creep up the sides of the funnel used for its filtration. Caution must be exercised whiletransferring the precipitate to the filter funnel to prevent its loss. Because DMG is a weakacid, the precipitate of nickel ion is somewhat sensitive to pH. Before precipitation, thesolution is buffered at basic pH with ammonia.

APPARATUS

Sintered glass filtering funnelSuction filtration apparatusOven (110 oC)Filter paperHot plateBeaker 600 mL

REAGENTS

Nickel sample pH paper1% dimethylglyoxime solution in alcohol6 M ammonia solutionTartaric acid

8/19/2019 Lab_Manual_STK1211_2015_2016

31/38

31

PROCEDURE

Part A: Preparation and Precipitation of the Nickel Sample

1. Clean a 600-mL beaker, rinse with distilled water.

2. Weigh out about 0.5 g of the unknown nickel sample into the beaker. Record theweight to the nearest mg (0.001 g).3. Add about 150 mL of distilled water to the beaker.4. Add approximately 0.2 g of tartaric acid. Stir to dissolve the solids. (The tartaric acid

forms soluble, stable metal complexes with any other metal ions that might be presentin a real sample and will prevent these other metals from precipitating with the nickelin this determination)

5. Adjust the pH of the sample to between 8 and 9, using 6 M ammonia solution drop bydrop. Do not dip pH paper into the sample. Rather, remove a single drop of thesolution with a stirring rod, and touch the drop to the pH paper.

6. Heat the sample on a hot plate to approximately 80 oC. After the sample has beenheated, remove the beaker from the hot plate.

7. Slowly and with constant stirring add 20 mL of 1% dimethylglyoxime solution. Anintensely red, fluffy precipitate should form at this point. If no precipitate forms,chances are the pH of your solution has not been correctly adjusted.

8. Allow the precipitate to settle until a layer of clear liquid is visible at the top of the beaker.

9. To ensure that precipitation is complete, add 2 – 3 drops of 1% DMG to the clearlayer. If no additional red precipitate forms, then you can assume that all of the nickelhas precipitated. If additional precipitate does form at this place, add 2 – 3 mLadditional 1% DMG and allow the precipitate to settle again. Then test the supernatant

liquid with another single drop of 1% DMG.10. Allow the sample beaker to cool to room temperature.

Part B: Filtering the Nickel/Dimethylglyoxime Precipitate

As it is being filtered, the Ni/DMG precipitate has a tendency of creeping up the sides ofthe filtering funnel. For this reason, never f il l the fil teri ng fun nel more than half ful l atany ti me.

1. Weigh a piece of filter paper.2. When the sample beaker have cooled to room temperature, set up filtering funnel into

the suction filtration apparatus. Fit the funnel with the weighed filter paper. Squirt thefilter paper with distilled water to moisten it.3. Turn on the suction, and slowly begin pouring the supernatant liquid from sample into

the funnel. When most of the liquid has been transferred, gradually begin transferringthe red precipitate. Remember not to fill the funnel more than half full at any time.

4. When the bulk of the precipitate has been transferred from the beaker to the funnel,use small portions of distilled water to transfer any remaining particles from the

beaker.5. When all the precipitate has been transferred to the funnel, increase suction to the

maximum level, and draw air through the precipitate for two minutes to help dry it.

8/19/2019 Lab_Manual_STK1211_2015_2016

32/38

32

6. Transfer the filter paper with the precipitate to a watch glass and dry the filter paperwith precipitate in the oven for at least one hour to remove all moisture. Make sure thatthe oven temperature is near 100 oC.

7. When the precipitates have dried completely, weigh the filter paper with the precipitate(to the nearest milligram).

8. Compare your data with two other groups in your laboratory.9. Calculate the mean % Ni in the original sample.

Part C: Calculations

The Ni/DMG precipitate is known to contain 20.32% nickel by weight.The mass of nickel present in a given amount of precipitate is then:

(mass of precipitate) (0.2032)

The percent nickel in the original sample is then:

100%takensampleof Mass

nickelof Mass

QUESTIONS

1. The percentage of nickel in the Ni/DMG precipitate is 20.32%. Using the atomic massof Ni and the formula mass of DMG, derive the percentage.

2. During the experiment, it was necessary to test the nickel sample for complete precipitation of nickel by adding a single drop of DMG to the supernatant solution.What error would be introduced into the percent Ni determined for a sample if not allthe nickel had been precipitated?

3. The DMG id provided as a solution in alcohol. Why is alcohol, rather than water, usedas the solvent?

8/19/2019 Lab_Manual_STK1211_2015_2016

33/38

33

Experiment 5: Gravimetric Determination of Nickel(II) Ion

Data Sheet

Name:

Student No.:

Date: _______________________________________________________________________

Particulars Sample 1 Sample 2 Sample 3

Mass of original sample (g)

Mass of filter paper (g)

Mass of filter paper with precipitate (g)

Mass of precipitate (g)

Mass of Ni in precipitate (g)

% Ni in original sample

Mean % Ni in original sample

8/19/2019 Lab_Manual_STK1211_2015_2016

34/38

34

EXPERIMENT 6: THIN-LAYER CHROMATOGRAPHY

INTRODUCTION

The word chromatography means colour-writing. The name was chosen when the methodwas first used to separate coloured components from plant leaves. Chromatography in itsvarious forms is perhaps the most important known method of chemical analysis ofmixtures.

Thin-layer chromatography (TLC) is a simple technique that can be used to separatemixtures into the individual components of the mixture. It uses a thin coating ofaluminium oxide (alumina) or silica gel on a glass slide or plaster sheet to which themixture to be resolved is applied. A single spot of the unknown mixture to be analyzed isapplied about 1 cm from the end of a strip of a TLC slide. The slide is then placed in ashallow layer of solvent or solvent mixture in a jar or beaker. Since the coating of the TLCslide is permeable to liquids, the solvent begins rising by capillary action.

As the solvent rises to the level at which the spot of mixture was applied, various effectscan occur, depending on the constituents of the spot. Those components of the spot thatare completely soluble in the solvent will be swept along with the solvent front as itcontinues to rise. Those components that are not at all soluble in the solvent will be left

behind at the original location of the spot. Most components of the unknown spot mixturewill take an intermediate approach as the solvent front passes. Components in the spot thatare somewhat soluble in the solvent will be swept along by the solvent front, but todifferent extents, reflecting their specific solubilities. By this means, the original spot ofmixture is spread out into a series of spots or band, with each spot representing one singlecomponent of the original mixture.

The separation of a mixture by chromatography is not solely a function of the solubility ofthe components in the solvent used. The TLC slide coating used in chromatography is notinert, but consists of molecules that may interact with the molecules of the components ofthe mixture being separated. Each component of the mixture is likely to have a differentextent of interaction with the slide coating. This differing extent of interaction between thecomponents of a mixture and the molecules of the support forms an equally important

basis for the separation. The TLC slide coating adsorbs molecules on its surface todiffering extents, depending on the structure and properties of the molecules involved.

To place a TLC separation on a quantitative basis, a mathematical function called theretention factor, Rf , is defined:

by solvent traveled distanceby spot traveled distance

R f

The retention factor depends on what solvent is used for the separation and on the specificcomposition of the slide coating used for a particular analysis. Because the retentionfactors for particular components of a mixture may vary if an analysis is repeated underdifferent conditions, a known sample is generally analyzed at the same time as anunknown mixture on the same slide. If the unknown mixture produces spots having the Rf values as spots from the known sample, then an identification of the unknown components

has been achieved.

8/19/2019 Lab_Manual_STK1211_2015_2016

35/38

35

Indicators are organic compounds that are typically used to signal a change in pH inacid/base titration analyses. Such indicators are dyes that exist in different coloured formsat different pHs, and change in colour of the indicator is the signal that the titrationanalysis is complete.

In this experiment, you will perform a thin-layer chromatographic analysis of a mixture ofthe dyes bromocresol green, methyl red, and xylenol orange. These dyes have been chosen

because they have significantly different retention factors, and a nearly completeseparation should be possible in the appropriate solvent system. You will also investigatethe effect of the solvent on TLC analyses, by attempting the separation in several differentsolvent systems.

APPARATUS

TLC slidesBeakersParafilmMicropipetteRulerPencil

REAGENTS

Methyl redXylenol orangeBromocresol greenAcetoneEthyl acetateHexaneEthanol

PROCEDURE

1. Clean and dry six beakers to be used as the chambers for the chromatography.2. Prepare mixtures of the solvents below, in the proportions indicated by volume, and

transfer each to a separate beaker. Cover the beakers with parafilm after adding thesolvent mixture, and label the beakers with the identity of the mixture it contains.

Acetone : hexane (3 : 2)Ethyl acetate : hexane (3 : 2)Acetone : ethyl acetate (1 : 1)Acetone : ethanol (1 : 1)Ethyl acetate : ethanol (1 : 1)Hexane : ethanol (1 : 1)

3. Prepare six TLC slides by marking lightly with pencil (not ink) a line across both thetop and bottom of the slide. Do not mark the line too deeply or you will remove the

coating of the slide. See Figure 6.1

8/19/2019 Lab_Manual_STK1211_2015_2016

36/38

36

R X G M

4. On one of the lines on each slide, mark four small pencil dots (to represent where thespots are to be applied). Above the other line on each slide, mark the following letters:R (methyl red), X (xylenol orange), G (bromocresol green) and M (mixture). SeeFigure 6.1.

5. Obtain small samples of the ethanol solutions of the three dyes and the mixture.Apply a single small droplet of the appropriate dye and the mixture to its pencil spoton each of the TLC slides. Use a separate micropipette for each dye and the mixture.Keep the spots of dye and the mixture as small as possible.

6. Allow the spots on the TLC slides to dry. Then gently lower one of the TLC slides,spots downward, into one of the solvent systems. Be careful not to wet the spots, or toslosh the solvent in the beaker. Do not move or disturb the beaker after adding theTLC slide. Carefully cover the beaker with parafilm.

7. Allow the solvent to rise on the TLC slide until it reaches the upper pencil line. Then

remove the TLC slide and quickly mark the exact solvent front before it evaporates.Mark the TLC slide with the identity of the solvent system used for development. Setthe TLC slide aside to dry completely.

8. Repeat the process using the additional TLC slides and solvent systems. Be certain tomark each slide with the solvent system used.

9. Determine Rf for each dye in each solvent system and record. Keep the TLC slidesand staple to the lab report for this experiment.

QUESTIONS

1. Why is it important to keep the spots applied to TLC slide for chromatography assmall as possible?

2. Why is it necessary to keep the beaker used for chromatography tightly covered with parafilm while the solvent is rising through the TLC slide?

3. Of the solvents used, some were very polar (eg. acetone, ethanol) while others werevery nonpolar (eg. hexane). Does the polarity of the solvent of the various solventmixtures seem to affect the completeness of the separation of dyes? Why might this

be so?

Figure 6.1

8/19/2019 Lab_Manual_STK1211_2015_2016

37/38

37

Experiment 6: Thin-layer Chromatography

Data Sheet

Name:

Student No.:

Date: _______________________________________________________________________

For each of the solvent mixtures studied, calculate Rf for each of the spot:

Acetone : hexane Distance travelled by solvent front _______________________

Distance travelled by spot Calculated Rf

Methyl red _____________________ _______________________

Xylenol orange _____________________ _______________________

Bromocresol green _____________________ _______________________

Ethyl acetate : hexane Distance travelled by solvent front _______________________


Methyl red _____________________ _______________________

Xylenol orange _____________________ _______________________


Acetone : ethyl acetate Distance travelled by solvent front _______________________


Methyl red _____________________ _______________________

Xylenol orange _____________________ _______________________


8/19/2019 Lab_Manual_STK1211_2015_2016

38/38

Acetone : ethanol Distance travelled by solvent front _______________________


Methyl red _____________________ _______________________

Xylenol orange _____________________ _______________________


Ethyl acetate : ethanol Distance travelled by solvent front _______________________


Methyl red _____________________ _______________________

Xylenol orange _____________________ _______________________


Hexane : ethanol Distance travelled by solvent front _______________________


Methyl red _____________________ _______________________

Xylenol orange _____________________ _______________________


Which solvent mixture gave the most complete resolution of the three dyes? Whichsolvent mixture gave the poorest resolution?

________________________________________________________________________

Lab_Manual_STK1211_2015_2016

Documents