University of Nizwa

Nizwa, Sultanate of Oman

CHEM 340ANALYTICAL CHEMISTRY FOR

PHARMACY LABORATORY MANUAL

Academic Year 2010/2011

By

DR. FAZAL MABOOD&

FRANCISCO C. CAMACHOHANAN ALYAHYAI

HAMEEDA ALRABAANIKHALID PERVEEZ

COLLEGE OF ARTS AND SCIENCES

DEPARTMENT OF BIOLOGICAL SCIENCES AND CHEMISTRY

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CHEM 340 – LABORATORY SYLLABUS

STUDY MATERIALS:

Douglas A. Skoog, Donald M. West, F. James Holler, Stanley R. Crouch, “Fundamentals of Analytical Chemistry” 8th ed. 2003; Saunders College Publishing, Philadelphia

TECHNIQUES INVOLVED:

This semester laboratory deals with the analytical techniques like volumetric techniques, spectroscopic techniques, and electroanalytical techniques used for the analysis of drug samples.

METHOD OF INSTRUCTION:

Demonstration will be given to the students in the Laboratory prior to start each experiment. Laboratory experiments will serve to reinforce, supplement, and provide a context for the materials presented in the theory class of course CHEM 340.

ASSESSMENT Total marks for lab exam 25

laboratory reports (5 marks ) Viva Voice exam (5 marks) For overall performance and practical lab exam (15 marks)

EXPECTED LEARNER OUTCOMES:

The students will be able to demonstrate a working knowledge of the proper use of analytical techniques for the analysis of different types of analytical and drug samples.

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SAFETY: The students must agree and sign the safety sheets before they start the first laboratory session.

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LABORATORY SAFETY AND GUIDELINES

Better Safe than Sorry!!!

Read the following instructions before you go into the laboratory

1. Safety glasses are required at all times. Prescription glasses are adequate, but contact lenses are not. Therefore, contact lenses are not permitted in the laboratory.

2. Always wear your laboratory coat. Do not wear clothing that hinders free movement of your hands or hangs loose outside your laboratory coat. Ladies Head scarves must be tucked into the lab coat.

3. Do not work in a laboratory if no instructor is present. Read theexperimental instructions carefully before starting the work. Note any precautionsthat must be taken.

4. Never eat, drink, or smoke in the laboratory. Never taste chemicals. Wash yourhands before you leave the laboratory.

5. Do not use your mouth to fill a pipette. There are special pipette fillers available.

6. Note the position of safety equipment such as fire extinguishers, eyewashes and first aid boxes. Report all accidents immediately to a staff member or technician.

7. Use the fume hood when doing experiments with irritatingchemicals.

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8. Be careful about throwing away wastes. Always follow instructions given. Do not throw away solids in the sink. Do not leave glassware or any other solid materials, including filter papers, in the sink. Put broken glassware into the labeled buckets. Some waste liquids must be put into special bottles, not poured down the sink.

9. Do not leave a lighted burner unattended. Do not allow any part of your body or of your clothing to fall into a flame.

10. When heating anything in a test tube, do not point the open end towards yourself or towards any other person.

11. Before you leave the laboratory, turn off any water taps and burners and dispose of solid waste in the correct container. Also wash out all apparatus used and clean up the bench top.

12. Keep your bench clean and tidy while you are working. Clean up any spills orbroken glass immediately. Keep your books and papers away from water, chemicalsand flames. Position your apparatus on the bench so that it is convenient andcomfortable to use. Put equipment you are not using out of the way, so that you donot knock it over. :

13. Finally If you are in any doubt ask.

I. General Behavior

1. Always read the upcoming experiment carefully and thoroughly before entering the laboratory.

2. Be in the lab and ready promptly when the lab begins.

3. Absolutely no food, beverages or chewing gums will be permitted 6

inside the lab.

4. Wash you hands frequently during the lab and of course wash them twice at the end of the lab.

5. Should any injury occur, regardless how minor it is, report it immediately to your supervisor.

6. Never pick up any broken glassware with your bare hands, regardless of the size of the pieces.

7. Never put broken glasses in a regular garbage can. A special container will be provided

8. Make sure to label and read the labels of all chemicals which your are dealing with.

9. Never use reagents from unmarked bottles.

10. In any emergency, the fastest to way to get the supervisor attention is to SCREAM.

If you are not feeling well, report it immediately to your supervisor

II. Safety Guidance and Tips

A. Do Not Pipette By Mouth - Ever

You say, "But it's only water." Even if it is, how clean do you think that glassware really is? Using disposable pipettes? I know lots of people who rinse them and put them back! Learn to use the pipette bulb or automated pipetter. Don't pipette by mouth at home either. Gasoline and kerosene should be obvious, but people get hospitalized or die every year, right? I know someone who used his mouth to start the suction on a waterbed to drain it. Do you know what they put in some waterbed additives?

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B. Identify the Safety Equipment

And know how to use it! Given that some people (possibly you) will need them, know the locations of the fire blanket, extinguishers, eyewash, and shower. Ask for demonstrations! If the eyewash hasn't been used in a while the discoloration of the water is usually sufficient to inspire use of safety glasses.

C. Don't Taste or Sniff Chemicals

For many chemicals, if you can smell them then you are exposing yourself to a dose that can harm you! If the safety information says that a chemical should only be used inside a fume hood, then don't use it anywhere else. This isn't cooking class - don't taste your experiments! D. Don't Casually Dispose of Chemicals Down the Drain

Some chemicals can be washed down the drain, while others require a different method of disposal. If a chemical can go in the sink, be sure to wash it away rather than risk an unexpected reaction between chemical 'leftovers' later.

E. Don't Eat or Drink in Lab

It's tempting, but oh so dangerous... just don't do it!

F. Don't Play Mad Scientist

Don't haphazardly mix chemicals! Pay attention to the order in which chemicals are to be added to each other and do not deviate from the instructions. Even chemicals that mix to produce seemingly safe products should be handled carefully. For example, hydrochloric acid and sodium hydroxide will give you salt water, but the reaction could

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break your glassware or splash the reactants onto you if you aren't careful!

G. Take Data During Lab

Not after lab, on the assumption that it will be neater. Put data directly in your lab book rather than transcribing from another source (e.g., notebook or lab partner). There are lots of reasons for this, but the practical one is that it is much harder for the data to get lost in your lab book. For some experiments, it may be helpful to take data before lab. No, I'm not telling you to dry-lab or cheat, but being able to project likely data will help you catch bad lab procedure before you are three hours or so into a project. Know what to expect. You should always read the experiment in advance. NOTE:

Failure to follow these rules may automatically and without warning result in a deduction of marks. In repeated cases you may be asked to leave the laboratory.

III. Hazard Control Information

On this drawing, please mark the locations of the following equipment. There may be more that one item of each kind. Please mark them all. Walk around the lab to find each item. You may need help from your instructor or a technician. Here is the list of items to find:

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1. All exit doors2. Fume hoods3. Fire extinguishers4. Eye wash stations

5. Fife blankets6. Emergency shower7. First aid kits8. Poster with safety instructions

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The Typical Experiment has three to five separate sections as follows:

• Pre-Laboratory Assignment: Involves calculations closely connected with the laboratory. This must be completed before you come to the lab.

•Introduction: Provides general interest and background information and a summary of the experiment.

• Experimental Procedure: You must read through this before you come to your laboratory session, paying particular attention to safety information.

• Calculations: Some laboratory experiments have an example of typical calculations to help you in carrying out the ones required for your data.

• Report form: You must record your data on the report form as you acquire it. Never write your data on scrap paper with the intention of transferring it later to the report form. Although this may lead to a neater report form, there is a great danger that you will lose the data. (If your data have been lost you cannot perform the necessary calculations. At best, you will have to repeat the laboratory experiment. At worst, you will be given a failing grade for that experiment.)

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NAME: ________________________________________BENCH NUMBER: ______________Instructor: __________________________________

I, the undersigned student, have received the safety training, understood it and agree to abide by the safety guidelines. I understand the importance of eye and body protection. If I therefore fail to abide by the safety rules; I am doing so at my own risk and will not hold University of Nizwa or Dr. Fazal Mabood liable for any injury that result.

Signature: ___________________________Date: _____________________

Student Name:_______________________________ID#:____________Date:_____________

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Table of Contents

General Laboratory InstructionsGeneral Laboratory ApparatusesLaboratory Safety and Guidelines

Exp. # 1Preparation and Standardization of 0.1M NaOH Solution With KHP by Using Acid Base Titration

Exp. # 2Preparation and Standardization of 0.1N HCl

Solution against 0.1N Na2CO3 Primary Standard Solution by Using Acid Base Titration

Exp. # 3Determination of %age of Acetic Acid in commercial Vinegar by Using Acid Base

Titration

Exp. # 4Preparation of a Buffer Solution of pH 5 and

then Determination of the effect of the Strong Acid on Buffer and Un Buffered Solution

Exp. # 5 Determination of Al +3 in Antacid Tablet (Trisil

Tablet) Using Acid Base Titration (Back Titration)

Exp. # 6Determination of Acetylsalicylic Acid in Aspirin

Tablet by Using Acid Base Titration

Exp. # 7Determination of % age Composition of Na2CO3 in Commercial Washing Soda Sample by Using

Acid Base Titration

Exp. # 8 Determination of the Contents of NaOH & Na2CO3 in the Given Mixture by Using Acid Base

Titration

Exp. # 9 Preparation & Standardization of 0.1N KMnO4 Solution using Na2C2O4 by Redox Titration

Exp. # 10 Determination of Fe2+ in a Given Solution by Redox Titration

Exp. # 11 Determination of % age Purity of Oxalic Acid by Redox Titration

Exp. # 12 Preparation and Standardization of 0.1M EDTA Solution with 0.1M Magnesium Solution by

Using Complexometric Titration

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Exp. # 13To Determine the Hardness of Water by

Complexometric Titration

Exp. # 14 Wavelength Optimization for Determination of KMNO4 by Using Spectrophotometric Method

Exp. # 15 Verification of Beers Law by Using Spectrophotometric Method

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EXPERIMENT 1Preparation and Standardization of 0.1M NaOH Solution with KHP by Using Acid Base Titration

Apparatus: Burette, pipette, funnel, beakers, flasks, wash bottle etcInstrument: BalanceReagents: NaOH, KHP & phenolphthalein indicatorChemical reaction:

TheoryStandardization is the process used for making to know the exact concentration of an unknown solution. In this process solutions of primary standard as well as of analyte are prepared. Then known volume from standard solution is taken in the volumetric flask and one drop of indicator solution is also added to this solution and then titrated against analyte solution that is taken in the burette. The end point will be the color change at that point volume used from burette should be measured and using the relationship of millimoles as,

Millimoles of analyte = Millimoles of primary standard solutionM1V1 = M2V2

Primary StandardsPrimary standards are those chemical reagents having high percent purity, stability toward air, high molecular mass, readily solubility in the solvent, medium cost and readily availability of the reagents. A few primary standards are as follow,

Potassium hydrogen phthalate (KHP) is used as a primary standard for standardization of NaOH base

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Na2CO3 is used as a primary standard for standardization of HCl acid

Na2C2O4 is used as a primary standard for standardization of KMnO4

Zn pellets or Mg ribbons are used as a primary standard for standardization of EDTA

Secondary Standard SolutionSecondary standard solution is that solution used for further standardization of solution having not exact known concentration. The preparation of secondary standard solution is done with the help of primary standard solution. After standardization with primary standard then we call it secondary standard solution and can be used for further standardization. For example the NaOH solution after standardization against KHP is now a secondary standard solution and this solution can be used for standardization of HCl solution. Procedure:To prepare 250ml volume of 0.1M NaOH solution, take 1g weight of NaOH with the help of balance in a watch glass, and then transfer this weight to a beaker and dissolve it in a few ml of distilled water and after dissolution transfer it into 250ml volumetric flask and dilute it upto the mark with distilled water. Now to standardize this solution against KHP solution, prepare 100ml volume of 0.1M KHP solution by dissolving 2.04g of KHP weight. Put NaOH solution into burette and transfer 10ml volume of KHP solution into titration flask and also add 1-drop of phenolphthalein indicator to this solution. After addition of indicator titrate this KHP solution against NaOH solution until get a pink color, that is the end point of titration and note the used volume of NaOH from burette. Take three readings of titration.Calculations for solutions preparation

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Two solutions we have to prepare1. Preparation of 0.1M NaOH solution of 250ml volume,To calculate weight for preparation of this solution we need to use formulaWeight in grams of NaOH =Molecular mass of NaOH x M x Volume in mL............(2) 1000M.mass of NaOH =23+16+1 =40amuSo put the values in above expression asWeight in grams of NaOH = 40 x0.1 M x 250 mL............(2) 1000Weight in grams of NaOH = 1gIt means to prepare 0.1M solution of NaOH of 250ml volume capacity we need to dissolve 1gram weight of NaOH that we calculated by using above expression 2. Preparation of 0.1M KHP solution of 100ml volumeAs we know the expressionWeight in grams of KHP =Molecular mass of KHP x M x Volume in mL............(2) 1000 M.mass of KHP = 204.2amuSo put the values in above expression asWeight in grams of KHP =204 x 0.1 x 100 mL............(2) 1000Weight in gram of KHP = 2.04gIt means to prepare 0.1M solution of KHP of 100ml volume capacity we need to dissolve 2.04gram weight of KHP and then dissolve in distilled water and after that dilute it upto the mark using 100ml volumetric flask After knowing the end point of titration use the following formula for making further calculations

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NaOH=KHPM1V1=M2V2

Observations & Calculations

S. No. Initial Volume (ml)

Final Volume (ml)

Difference Volume (ml)

123

Mean Volume =

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Calculations

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EXPERIMENT 2Preparation and Standardization of 0.1N HCl Solution

against 0.1N Na2CO3 Primary Standard Solution by Using Acid Base Titration

Apparatus: Burette, pipette, funnel, beakers, flasks, wash bottle etcInstrument: BalanceReagents: HCl, Na2CO3 & Methyl orange indicatorChemical reaction:

TheoryStandardization is the process used for making to know the exact concentration of an unknown solution. In this process solutions of primary standard as well as of analyte are prepared. Then known volume from standard solution is taken in the volumetric flask and one drop of indicator solution is also added to this solution and then titrated against analyte solution that is taken in the burette. The end point will be the color change at that point volume used from burette should be measured and using the relationship of millimoles as,

Millimoles of analyte = Millimoles of primary standard solutionM1V1 = M2V2

Procedure:To prepare 250ml solution of 0.1N HCl, take 2.05 ml volume of HCl from reagent bottle and dilute upto the mark with distilled water using 250ml volumetric flask. Now to standardize this solution against Na2CO3 primary standard solution, prepare 100ml volume of 0.1N Na2CO3 solution by dissolving 0.53g of Na2CO3 weight. Put HCl solution into burette and transfer 10ml volume of Na2CO3 solution into titration flask and also add 1-drop of Methyl orange indicator to this solution.

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After addition of indicator titrate this Na2CO3 solution against HCl solution until get a pink color, that is the end point of titration and note the used volume of HCl from burette. Take three readings of titration.Calculations for solutions preparationTwo solutions we have to prepare

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Preparation of 0.1N HCl solution of 250ml volume,To calculate weight for preparation of this solution we need to use formulaWeight in grams of HCl =Eq. mass of HCl xN x Volume in mL............(2) 1000Eq.mass of HCl =36.5amuSo put the values in above expression asWeight in grams of HCl = 36.5 x0.1 N x 250 mL............(2) 1000Weight in grams of NaOH = 0.912gAs HCl is liquid so will be taking volume instead of mass using density mass relationship asD = m, VOr V= m DAs density of HCl is 1.19g/ml, so put the values in above expression asV= 0.912 =0.76ml of pure HCl 1.19But the available HCl reagent bottle HCl having 37% purity it means

= 2.07ml of reagent bottle HClIt means to prepare 0.1N solution of HCl of 250ml volume capacity we need to dissolve 2.07ml of HCl that we calculated by using above expression Preparation of 0.1N Na2CO3 solution of 100ml volumeAs we know the expression

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Weight in grams of Na2CO3 =Molecular mass of Na2CO3 x N x Volume in mL............(2) 1000Eq.mass of Na2CO3 = 53amuSo put the values in above expression asWeight in grams of Na2CO3 =53 x 0.1 x 100 mL............(2) 1000Weight in gram of KHP = 0.53gIt means to prepare 0.1N solution of Na2CO3 of 100ml volume capacity we need to dissolve 0.53gram weight of Na2CO3 and then dissolve in distilled water and after that dilute it upto the mark of 100ml volumetric flask.After knowing the end point of titration use the following formula for making further calculations

HCl= Na2CO3

N1V1==N2V2Observations & Calculations

S. No. Initial Volume (ml)

Final Volume (ml)

Difference Volume (ml)

123

Mean Volume =

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EXPERIMENT 3Determination of %age of Acetic Acid in commercial

Vinegar by Using Acid Base TitrationApparatus: Burette, pipette, funnel, beakers, flasks, wash bottle etcInstrument: BalanceReagents: NaOH,, Vinegar sample, KHP & Phenolphthalein indicatorChemical reaction:

Theory

Titration is an analytical technique used to determine the concentration of an unknown substance. Titration involves the slow addition of one solution where the concentration is known called standard solution to a known volume of another solution where the concentration is unknown called analyte until the reaction reaches a desired level. The point at which all of the analyte is consumed is called the endpoint and is determined by some type of indicator that is also present in the solution.

Acid-base titrations are based on the on acid base reactions and are also called neutralization titration.Acid-base titrations are used to determine the unknown amount of most acids and bases. There are many standard substances that can be used in acid base titrations. Those most popular are sodium carbonate Na2CO3, borax (disodium tetraborate decahydrate) Na2B4O7·10H2O and potassium hydrogen phthalate KHC8H4O4, often called simply KHP.Type of indicator depends on several factors. One of them is the equivalence point pH. Depending on the titrated substance and titrant used this can vary, usually between 4 and 10. However, even if it is often possible (see list of pH indicators) we are rarely selecting

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indicator that changes color exactly at the equivalence point, as usually increase of accuracy doesn't justify additional costs. Thus in practice you will probably use phenolphtalein when NaOH is used as the titrant and methyl orange when titrating with the strong acid.

ProcedurePrepare 250 ml solution of 0.1N NaOH and standardize it against 0.1N KHP solution. Take 2ml of vinegar sample in a titration flask add 1 drop of phenolphthalein indicator to this vinegar sample and titrate it against standard solution of NaOH taken in burette. The end point will be appearance of pink color. Note the volume of NaOH solution used for color change.Observations & Calculations

S. No. Initial Volume (ml)

Final Volume (ml)

Difference Volume (ml)

123

Mean Volume =

26

EXPERIMENT 4Preparation of a Buffer Solution of pH 5 and then Determination of the effect of the Strong Acid on

Buffer and Un Buffered SolutionTheoryBuffer solutions are those solutions when they are added to some other solution they resist to change in pH of that solution if small amount of strong acid or strong base is added to that solution. Or We can say the buffer solution has the property to control the pH of that solution to which buffer solution has been added.Buffer solution are normally made of either weak acid or weak base and the salt derived from that weak acid, or weak base e.g. buffer solution of CH3COOH/CH3COONa and the other one NH3OH/NH4Cl Solutions Preparationa) 0.1N CH3COONa Solution of 100ml volume

Dissolve 0.82g of CH3COONa salt in distilled water and dilute it upto the mark with distilled water using 100ml volumetric flask.b) 0.1N CH3COOH Solutions of 100ml volume

Dissolve 0.57ml of CH3COOH in distilled water and dilute it upto the mark with distilled water using 100ml volumetric flask.c) Buffer solution of pH5To prepare a buffer solution of pH5 mix 65ml of 0.1N acetic acid solution and 35ml of 0.1N sodium acetate solution in a 100ml volumetric flask.d) 0.01N HCl Solution of 100ml volume

Dissolve 1ml of HCl in distilled water and dilute it upto the mark with distilled water using 100ml volumetric flask.

Procedure

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Transfer buffer solution to one beaker and take 100ml of distilled water in other beaker. Then insert the glass electrode of pH meter in both of the beakers solutions and record the pH of both of solutions. After that add 2ml of 0.01N HCl solution into both of the beakers and again record their pH.

Name of solution

pH before addition of HCl solution

pH after addition of HCl solution

Difference in pH

Buffer solution

Distilled water

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EXPERIMENT 5Determination of Al +3 in Antacid Tablet (Trisil Tablet)

Using Acid Base Titration (Back Titration)Apparatus: Burette, pipette, funnel, beakers, flasks, wash bottle etcInstrument: BalanceReagents: HCl, NaOH, Trisil tablet, KHP & Phenolphthalein indicator

Chemical reaction:

Theory

Sometimes it is not possible to use standard titration methods. For example the reaction between determined substance and titrant can be too slow, or there can be a problem with end point determination. In such situations we can often use a technique called back titration. In back titration we use two reagents - one, that reacts with the original sample (lets call it A), and second (lets call it B), that reacts with the first reagent. How do we proceed? We add precisely measured amount of reagent A to sample and once the reaction ends we titrate excess reagent A left with reagent B. Knowing initial amount of reagent A and amount that was left after the reaction (from titration) we can easily calculate how much reagent A was used for the first reaction.

Procedure:Prepare 0.1N solution of NaOH & 0.1N HCl each of 250ml volume capacity and then standardized both of those solutions using primary standard solutions. Take one Trisil tablet, weight it and then dissolve it in distilled water and dilute upto to the mark using 100ml volumetric

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flask. Take 20ml volume from the Trisil tablet sample solution in a titration flask and to this solution add 40ml of 0.1N HCl standard solution and also add one drop of phenolphthalein indicator. Titrate this mixture against standard solution of NaOH solution. End point of titration will be appearance of pink color.

Observations & Calculations

S. No. Initial Volume (ml)

Final Volume (ml)

Difference Volume (ml)

123

Mean Volume =

30

EXPERIMENT 6

Determination of Acetylsalicylic Acid in Aspirin Tablet by Using Acid Base Titration

Apparatus: Burette, pipette, funnel, beakers, flasks, wash bottle etcInstrument: BalanceReagents: NaOH, Aspirin tablet & Phenolphthalein indicator

Chemical reaction:

Theory

Titration is an analytical technique used to determine the concentration of an unknown substance. Titration involves the slow addition of one solution where the concentration is known called standard solution to a known volume of another solution where the concentration is unknown called analyte until the reaction reaches a desired level. The point at which all of the analyte is consumed is called the endpoint and is determined by some type of indicator that is also present in the solution.

Acid-base titrations are based on the on acid base reactions and are also called neutralization titration.Acid-base titrations are used to determine the unknown amount of most acids and bases. There are many standard substances that can be used in acid base titrations. Those most popular are sodium carbonate Na2CO3, borax

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(disodium tetraborate decahydrate) Na2B4O7·10H2O and potassium hydrogen phthalate KHC8H4O4, often called simply KHP.Type of indicator depends on several factors. One of them is the equivalence point pH. Depending on the titrated substance and titrant used this can vary, usually between 4 and 10. However, even if it is often possible (see list of pH indicators) we are rarely selecting indicator that changes color exactly at the equivalence point, as usually increase of accuracy doesn't justify additional costs. Thus in practice you will probably use phenolphtalein when NaOH is used as the titrant and methyl orange when titrating with the strong acid.

Procedure:Prepare 0.01N solution of NaOH of 250ml volume capacity and then standardized it using KHP primary standard solution. Take 3 tablets of aspirin, weight it and then dissolve it in distilled water and dilute upto to the mark using 100ml volumetric flask. Take 20ml volume from the aspirin tablet sample solution in a titration flask and also add one drop of phenolphthalein indicator. Titrate this mixture against standard solution of NaOH solution. End point of titration will be appearance of pink color.Observations & Calculations

S. No. Initial Volume (ml)

Final Volume (ml)

Difference Volume (ml)

123

Mean Volume =

32

EXPERIMENT 7

Determination of % age Composition of Na2CO3 in Commercial Washing Soda Sample by Using Acid Base

TitrationApparatus: Burette, pipette, funnel, beakers, flasks, wash bottle etcInstrument: BalanceReagents: NaOH, Commercial washing soda, HCl, Methyl orange indicator

Chemical reaction:

Theory

Titration is an analytical technique used to determine the concentration of an unknown substance. Titration involves the slow addition of one solution where the concentration is known called standard solution to a known volume of another solution where the concentration is unknown called analyte until the reaction reaches a desired level. The point at which all of the analyte is consumed is called the endpoint and is determined by some type of indicator that is also present in the solution.

Acid-base titrations are based on the on acid base reactions and are also called neutralization titration.Acid-base titrations are used to determine the unknown amount of most acids and bases.

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There are many standard substances that can be used in acid base titrations. Those most popular are sodium carbonate Na2CO3, borax (disodium tetraborate decahydrate) Na2B4O7·10H2O and potassium hydrogen phthalate KHC8H4O4, often called simply KHP.Type of indicator depends on several factors. One of them is the equivalence point pH. Depending on the titrated substance and titrant used this can vary, usually between 4 and 10. However, even if it is often possible (see list of pH indicators) we are rarely selecting indicator that changes color exactly at the equivalence point, as usually increase of accuracy doesn't justify additional costs. Thus in practice you will probably use phenolphtalein when NaOH is used as the titrant and methyl orange when titrating with the strong acid.

Procedure:Prepare 0.01N solution of HCl of 250ml volume capacity and then standardized it using Na2CO3 primary standard solution. Take known weight (0.13g) of washing soda sample dissolve it in distilled water and transfer it into a titration flask and also add one drop of methyl orange indicator. Titrate this mixture against standard solution of HCl solution. End point of titration will be appearance of pink color.Repeat this procedure for three times to take mean at end point and calculate the weight of Na2CO3 by using the formula

Moles of Na2CO3 = Moles of HCl

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EXPERIMENT 8

Determination of the Contents of NaOH & Na2CO3 in the Given Mixture by Using Acid Base Titration

Apparatus: Burette, pipette, funnel, beakers, flasks, wash bottle etcInstrument: BalanceReagents: Mixture of NaOH & Na2CO3 , HCl, Methyl orange & phenolphthalein indicatorsChemical reaction:

Procedure:Prepare 0.01N solution of HCl of 250ml volume capacity and then standardized it against Na2CO3 primary standard solution. Take 10ml volume of mixture solution and transfer it into a titration flask and also add one drop of phenolphthalein indicator. Titrate this mixture against standard solution of HCl solution. End point of titration will be disappearance of pink color. Record the volume used at this point. This is the volume of HCl used for neutralization of NaOH & one hydroxyl of sodium carbonate. Now add a drop of methyl orange indicator to it and titrate it against standard solution of HCl until appearance of pink color. Again note the volume used for color change. Now subtract the first volume from the second volume. We will get the volume of HCl used for neutralization of NaHCO3. By subtracting volume of HCl for NaHCO3 from the first reading we will get the volume of HCl used for neutralization of NaOH. Weight of NaOH & Na2CO3 was calculated by using formula

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Observations & Calculations

S. No. Initial Volume (ml)

Final Volume (ml)

Difference Volume (ml)

123

Mean Volume =

37

EXPERIMENT 9

Preparation & Standardization of 0.1N KMnO4 Solution using Na2C2O4 by Redox Titration

Apparatus: Burette, pipette, funnel, beakers, flasks, wash bottle etcInstrument: Balance, Reagents: KMnO4, Na2C2O4, H2SO4

Chemical reaction:

Theory

Redox titrations are based on redox reactions. Redox reaction is that reaction in that one specie loss electron and the other specie can gain electron e.g.

KMnO4 + 5Fe2+ + H+ → Mn2+ + 5Fe3+

In the above reaction Mn7+ is reduced to Mn2+ by gaining five electrons from five atoms of iron and Fe2+ is oxidized to Fe3+.

There are many redox reagents used in redox titrations. To list a few - potassium permanganate is used for determination of Fe2+, H2O2 and oxalic acid. Potassium dichromate for determination of Fe2+ and Cu in CuCl. Bromate is used for tin and phenol, iodides (titrated with sodium thiosulfate) for H2O2 and Cu2+. Cerium (IV) can be used to determine ferrocyanides and nitrites. There are also many other methods.

Redox indicators

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In order to locate the end point in case of redox titration there are three types of indicators used i.e. true redox indicators, self indicators & specific indicators

Commonly used true redox indicators are substances that can exist in two forms i.e. oxidized and reduced form that differ in color. Potential at which the substance changes color must be such that the change occurs close to the equivalence point. Examples of such substances are ferroin, diphenylamine or nile blue

In the case of one color indicators, potential at which indicator color starts to be visible depends on the indicator concentration. The most obvious one is while the general idea that observed color depends on the ratio of concentrations of both reduced and oxidized forms still holds, ratio of concentrations is not pH dependent, but redox potential dependent. We can easily calculate ratio of the concentrations of both forms using Nernst equation:

Procedure:To prepare 0.1N KMnO4 solution of 250ml volume capacity, dissolve 0.79g weight of KMnO4 and dilute it upto 250ml volume using volumetric flask with distilled water. This weight was calculated using the formula as given below

Eq. mass of KMnO4 = 31.6Put KMnO4 solution into burette.Similarly to prepare 0.1N Na2C2O4solution of 100ml volume capacity, dissolve 0.67g weight of Na2C2O4 and dilute it upto 100ml volume using volumetric flask with distilled water.

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Take 10ml of 0.1N Na2C2O4 solution into a titration flask and also add 2ml of conc H2SO4 to this flask and titrate it against KMnO4 solution taken in burette. The end point of this titration will be appearance of pink color. Take three reading of this titration. Normality of KMnO4

solution can be calculated using this expression as belowKMnO4 = Na2C2O4

N1V1= N2V2

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Observations & Calculations

S. No. Initial Volume (ml)

Final Volume (ml)

Difference Volume (ml)

123

Mean Volume =

41

EXPERIMENT 10Determination of Fe2+ in a Given Solution by Redox

TitrationApparatus: Burette, pipette, funnel, beakers, flasks, wash bottle etcInstrument: Balance, Reagents: KMnO4, Na2C2O4, H2SO4, Unknown solutionChemical reaction:

Theory

Redox titrations are based on redox reactions. Redox reaction is that reaction in that one specie loss electron and the other specie can gain electron e.g.

KMnO4 + 5Fe2+ + H+ → Mn2+ + 5Fe3+

In the above reaction Mn7+ is reduced to Mn2+ by gaining five electrons from five atoms of iron and Fe2+ is oxidized to Fe3+.

There are many redox reagents used in redox titrations. To list a few - potassium permanganate is used for determination of Fe2+, H2O2 and oxalic acid. Potassium dichromate for determination of Fe2+ and Cu in CuCl. Bromate is used for tin and phenol, iodides (titrated with sodium thiosulfate) for H2O2 and Cu2+. Cerium (IV) can be used to determine ferrocyanides and nitrites. There are also many other methods.

Redox indicators

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In order to locate the end point in case of redox titration there are three types of indicators used i.e. true redox indicators, self indicators & specific indicators

Commonly used true redox indicators are substances that can exist in two forms i.e. oxidized and reduced form that differ in color. Potential at which the substance changes color must be such that the change occurs close to the equivalence point. Examples of such substances are ferroin, diphenylamine or nile blue

In the case of one color indicators, potential at which indicator color starts to be visible depends on the indicator concentration. The most obvious one is while the general idea that observed color depends on the ratio of concentrations of both reduced and oxidized forms still holds, ratio of concentrations is not pH dependent, but redox potential dependent. We can easily calculate ratio of the concentrations of both forms using Nernst equation:Procedure:Put standard KMnO4 solution into burette. Take 10ml of unknown

solution and 2ml of conc. H2SO4 and transfer it into a titration flask and titrate it against standard solution of KMnO4 solution taken in burette. The end point of this titration will be appearance of pink color. Take three reading of this titration. Weight in grams of iron can be calculated using this expression as below

Fe2+ = KMnO4

Observations & Calculations

S. No. Initial Volume Final Volume Difference Volume

43

(ml) (ml) (ml)123

Mean Volume =

44

EXPERIMENT 11Determination of % age Purity of Oxalic Acid by Redox

TitrationApparatus: Burette, pipette, funnel, beakers, flasks, wash bottle etcInstrument: Balance, Reagents: KMnO4, Na2C2O4, H2SO4, H2C2O4

Chemical reaction:

Theory

Redox titrations are based on redox reactions. Redox reaction is that reaction in that one specie loss electron and the other specie can gain electron e.g.

KMnO4 + 5Fe2+ + H+ → Mn2+ + 5Fe3+

In the above reaction Mn7+ is reduced to Mn2+ by gaining five electrons from five atoms of iron and Fe2+ is oxidized to Fe3+.

There are many redox reagents used in redox titrations. To list a few - potassium permanganate is used for determination of Fe2+, H2O2 and oxalic acid. Potassium dichromate for determination of Fe2+ and Cu in CuCl. Bromate is used for tin and phenol, iodides (titrated with sodium thiosulfate) for H2O2 and Cu2+. Cerium (IV) can be used to determine ferrocyanides and nitrites. There are also many other methods.

Redox indicators

45

In order to locate the end point in case of redox titration there are three types of indicators used i.e. true redox indicators, self indicators & specific indicators

Commonly used true redox indicators are substances that can exist in two forms i.e. oxidized and reduced form that differ in color. Potential at which the substance changes color must be such that the change occurs close to the equivalence point. Examples of such substances are ferroin, diphenylamine or nile blue

In the case of one color indicators, potential at which indicator color starts to be visible depends on the indicator concentration. The most obvious one is while the general idea that observed color depends on the ratio of concentrations of both reduced and oxidized forms still holds, ratio of concentrations is not pH dependent, but redox potential dependent. We can easily calculate ratio of the concentrations of both forms using Nernst equation:Procedure:Put standard solution KMnO4 solution into burette.Take 10ml of oxalic acid solution prepared by dissolving 0.45g of oxalic acid in 100ml volume and 2ml of conc H2SO4 and transfer it into a titration flask and titrate it against standard solution of KMnO4 solution taken in burette. The end point of this titration will be appearance of pink color. Take three reading of this titration. Weight in grams of iron can be calculated using this expression as below

Oxalic acid = KMnO4

Observations & Calculations

S. No. Initial Volume Final Volume Difference Volume

46

(ml) (ml) (ml)123

Mean Volume =

47

EXPERIMENT 12Preparation and Standardization of 0.1M EDTA

Solution with 0.1M Magnesium Solution by Using Complexometric Titration

Apparatus: Burette, pipette, funnel, beakers, flasks, wash bottle etcInstrument: Balance, Reagents: EDTA solution, magnesium ribbon, buffer solution of pH 10Indicator: Erichrome black T (EBT)Chemical reaction:

Theory

Complexometric titrations are based on complexation reactions. Most often complexating reagent is EDTA (EthyleneDiamineTetraAcetic acid). There are also other similar chelating agents (EGTA, CDTA and so on) used.

In the case of determination of metals detection of the endpoint is mainly based on substances that change color when creating complexes with determined metals. One of these indicators is Eriochrome black T, which is a weak complexating reagent complex with metal, other examples are pyrocatechin violet and murexide. It is important that formation constant for these complexes is low enough, so that titrant reacts with complexed ions first.

ProcedurePrepare a buffer solution of pH 10 by dissolving 7.1g of NH4Cl and 58.5ml of NH4OH in a 100ml volumetric flask and diluting upto the mark with distilled water.

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Similarly to prepare 0.1M solution of EDTA, dissolve 9.3g of sodium EDTA salt in a little of distilled water and dilute upto the mark using 250ml volumetric flask with distilled water.Similarly to prepare 0.1M solution of magnesium dissolve 1g of washed magnesium ribbon in a little of conc. HCl and then dilute it with distilled water upto the mark using 100ml volumetric flask.Now transfer 10ml of magnesium solution into a titration flask also add 5ml of pH 10 buffer solution to this flask and also add 1 drop of freshly prepared solution of EBT indicator. Titrate it against EDTA solution until the appearance of light blue color. It is the end point of titration.

EDTA solution= Mg solutionM1V1=M2V2

Observations & Calculations

S. No. Initial Volume (ml)

Final Volume (ml)

Difference Volume (ml)

123

Mean Volume =

49

EXPERIMENT 13

To Determine the Hardness of Water by Complexometric Titration

Apparatus: Burette, pipette, funnel, beakers, flasks, wash bottle etcInstrument: Balance, Reagents: EDTA solution, magnesium ribbon, buffer solution of pH 10Indicator: Erichrome black T (EBT)Chemical reaction:

Theory

Complexometric titrations are based on complexation reactions. Most often complexating reagent is EDTA (EthyleneDiamineTetraAcetic acid). There are also other similar chelating agents (EGTA, CDTA and so on) used.

In the case of determination of metals detection of the endpoint is mainly based on substances that change color when creating complexes with determined metals. One of these indicators is Eriochrome black T, which is a weak complexating reagent complex with metal, other examples are pyrocatechin violet and murexide. It is important that formation constant for these complexes is low enough, so that titrant reacts with complexed ions first.

ProcedurePut standard solution of EDTA in burette. Now take 10ml of water into a titration flask also add 5ml of pH 10 buffer solution to this flask and also add 1 drop of freshly prepared solution of EBT indicator. Titrate it

50

against EDTA solution until the appearance of light blue color. It is the end point of titration.

EDTA solution= Mg solutionM1V1=M2V2

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Observations & Calculations

S. No. Initial Volume (ml)

Final Volume (ml)

Difference Volume (ml)

123

Mean Volume =

52

EXPERIMENT 14

Wavelength Optimization for Determination of KMNO4 by Using Spectrophotometric Method

Apparatus: Burette, pipette, funnel, beakers, flasks, wash bottle etcInstrument: Balance, SpectrophotometerReagents: 100ppm, 2ppm, 4ppm, 6ppm, 8ppm & 10ppm KMnO4 solutions of 100ml volume capacity,

TheorySpectroscopy is a kind of analytical techniques used for determination of analyte. It is based on interaction of electromagnetic radiations with matter. Some time the matter absorbs or sometimes emits radiations. So in spectroscopy we measure the magnitude of absorbed or emitted radiations and then relate that magnitude of radiations with the concentration of analyte species.Spectroscopy is broadly categories into two classes i.e. Absorption Spectroscopy and Emission Spectroscopy. In case of absorption spectroscopy we pass electromagnetic radiations from analyte species and then measure the magnitude of radiation absorbed by analyte species and then relating the amount of absorbed radiation with the concentration of analyte species using Beer’s Lambert Law i.e.

A = bC

Where & b are constant so, A C (absorption is directly proportion to

concentration of analyte specie)

While in case of emission spectroscopy we measure the intensity of emitted

radiations and then relating the magnitude of intensity of emitted radiation to the

concentration of analyte species i.e. I C (Intensity is directly proportion to

concentration of analyte specie)

While in case of emission spectroscopy we measure the intensity of emitted radiations and then relating the magnitude of intensity of

53

emitted radiation to the concentration of analyte species i.e I C

(Intensity is directly proportion to concentration of analyte specie)

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Procedure:To prepare 100ppm KMnO4 of 100ml volume capacity, dissolve 0.01g weight of KMnO4 and dilute it upto 100ml volume using volumetric flask with distilled water. Prepare dilute solutions of 2ppm, 4ppm, 6ppm, 8ppm & 10ppm from the stock solution by taking 2ml, 4ml, 6ml, 8ml & 10ml volume from the stock solution and diluting to 100ml volume with distilled water using 100ml volumetric flasks. In order to optimize the wavelength take some volume of 6ppm solution in a covet and put it in a sample holder cabin of spectrophotometer. Before measuring the absorbance of this solution first set the transmittance with blank at 100% transmittance. After this put the solution covet in the instrument and measure its absorbance on 400nm wavelength. Now change the wavelength to 420nm and again repeating the same procedure measure the absorbance. Using the same procedure record the absorbance after change in wavelength of 20nm difference till to 600nm. Plot the absorbance against wavelength and find the wavelength at which there is maximum absorbance and that is called optimum wavelength.

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EXPERIMENT 15

Verification of Beers Law by Using Spectrophotometric Method

Apparatus: Burette, pipette, funnel, beakers, flasks, wash bottle etcInstrument: Balance, SpectrophotometerReagents: 100ppm, 2ppm, 4ppm, 6ppm, 8ppm & 10ppm KMnO4 solutions of 100ml volume capacity,

TheorySpectroscopy is a kind of analytical techniques used for determination of analyte. It is based on interaction of electromagnetic radiations with matter. Some time the matter absorbs or sometimes emits radiations. So in spectroscopy we measure the magnitude of absorbed or emitted radiations and then relate that magnitude of radiations with the concentration of analyte species.Spectroscopy is broadly categories into two classes i.e. Absorption Spectroscopy and Emission Spectroscopy. In case of absorption spectroscopy we pass electromagnetic radiations from analyte species and then measure the magnitude of radiation absorbed by analyte species and then relating the amount of absorbed radiation with the concentration of analyte species using Beer’s Lambert Law i.e.

A = bC

Where & b are constant so, A C (absorption is directly proportion to

concentration of analyte specie)

While in case of emission spectroscopy we measure the intensity of emitted

radiations and then relating the magnitude of intensity of emitted radiation to the

concentration of analyte species i.e. I C (Intensity is directly proportion to

concentration of analyte specie)

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While in case of emission spectroscopy we measure the intensity of emitted radiations and then relating the magnitude of intensity of emitted radiation to the concentration of analyte species i.e I C

(Intensity is directly proportion to concentration of analyte specie)

Procedure:To prepare 100ppm KMnO4 of 100ml volume capacity, dissolve 0.01g weight of KMnO4 and dilute it upto 100ml volume using volumetric flask with distilled water. To prepare standard solutions of 2ppm, 4ppm, 6ppm, 8ppm & 10ppm from the stock solution by taking 2ml, 4ml, 6ml, 8ml & 10ml volume from the stock solution and diluting to 100ml volume with distilled water using volumetric flasks. Now measure the absorbance of each standard solution i.e. 2ppm, 4ppm, 6ppm, 8ppm & 10ppm as well as sample solution at optimum wavelength i.e. 510nm and plot the absorbance against concentration i.e. standard calibration curve. With the help of this curve find the concentration of sample solution connecting the absorbance with concentration on the curve.

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