inorganic chemistry practical protocal

8/4/2019 Inorganic Chemistry Practical Protocal

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TABLE OF CONTENTS

1. Acid Base Titrations (Acidimetry and Alkalimetry)

2. Estimation of Hardness of Water

3. Estimation of Chloride in Water

4. Determination of COD of Industrial Waste Water

5. Estimation of CaO in Lime

6. Analysis of Haematite Ore Volumetrically

7. Analysis of Haematite Ore Gravimetrically

8. Analysis of Brass Volumetrically

9. Analysis of Brass Gravimetrically

10. Analysis of Stainless Steel


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ACID BASE TITRATIONS (ACIDIMETRY AND ALKALIMETRY)

Acidimetry

INTRODUCTION:

Volumetric analysis is a quantitative chemical analysis method that is used to determine

the volume of a solution of accurately known concentration (standard solution) which is required

to react quantitatively with a measured volume of a solution of unknown concentration. The

process of determination of volume of standard solution required to react completely with a

solution of the substance to be determined is called titration. The solution of known

concentration is called the titrant and the solution of known volume but unknown concentration

is called the analyte. The equivalence point is the point in a titration where stoichiometrically

equivalent amounts of analyte and titrant are present whereas the end point is the point at whichan observable physical change like color signals the equivalence point. The advantages of such

an analysis include high precision, require simpler apparatus and can be quickly performed.

A visual indicator is a compound that exhibits colors depending on the pH of its

surroundings. There are four types of titrations based on the reactions employed in volumetric

analysis; i) acid-base titrations (ii) redox titrations (iii) precipitation titrations and (iv)

complexometric titrations. The particular purpose of this experiment is to determine the strength

of given sulphuric acid solution using sodium carbonate crystals.

LEARNING OBJECTIVES:

a) To demonstrate the basic principles of titrations

b) To prepare standard solution

c) To determine the strength of the given sulphuric acid solution using sodium carbonate

crystals.

d) The use of indicators

PRINCIPLE:

A standard solution of sodium carbonate is prepared by dissolving a known weight of it in a

known volume of water. Then a known volume of this solution is titrated against the given

sulphuric acid solution taken in the burette. Since sulphuric acid is a strong acid and sodiumcarbonate is a weak base, methyl orange is used as the indicator.

MATERIALS AND METHODS:

anhydrous sodium carbonate, sulphuric acid solution, methyl orange indicator


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a. Preparation of standard Na2CO3 solution:

1. About 1.3g of pure anhydrous sodium carbonate is weighed accurately into a 250 mL standard flask.

2. It is dissolved in a little distilled water and then the solution is made up to the mark with distilled

water and shaken well for uniform concentration.

b. Standardisation of H2SO4 solution

1. The burette is filled with the given sulphuric acid solution.

2. Then 25 mL of the sodium carbonate solution is pipetted into the conical flask.

3. Three drops of methyl orange indicator is added to it and then the solution turns yellow.

4. The initial burette reading is noted.

5. The solution in the conical flask is then titrated against the sulphuric acid solution in the burette,

carefully shaking the solution all the times.

6. Towards the end of the titration, solution from the burette is added dropwise.

7. Change of colour from yellow to orange red marks the end point.

8. The burette reading is again noted.

9. The titration is repeated to get the concordant values.

DISCUSSION:

Acidimetry is a simple experiment, which is useful and informative in demonstrating the basic principlesof titrations. The importance of such simple experiments cannot be overestimated, as they directly

represent the identity and image of a quantitative analysis and serve as evidence of scientific value

facilitating the technical skill upgradation. The concentration of analyte is determined from the

stoichiometry of the reaction and the volume of titrant required to carry it out.

Result:

The strength of the given sulphuric acid solution =

Observations and Calculation

Preparation of standard Na2CO3 solution

Weight of the weighing bottle + Na2CO3 crystals, W1 =

Weight of the empty weighing bottle, W2 =

Weight of Na2CO3 crystals taken, ( W1 - W2) =

Strength ofNa2CO3 solution prepared = ( W1 - W2)*4 / 53

=


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=

Standardisation of sulphuric acid

Solution taken in the burette given H2SO4 solution

Solution taken in the flask 25 mL of Na2CO

3solution

Indicator used 3 drops ofmethyl orange

Colour change yellow to orange red

Trial No. 1 2 3

Final Burette reading

Initial Burette reading

Volume of acid solution added

Agreeing value, V =

Therefore, Normality of the given H2SO4 solution,

N H2SO4 = (V Na2CO3 * N Na2CO3) / V H2SO4

=

=

Alkalimetry

PRINCIPLE:

A standard solution of oxalic acid is prepared by dissolving a known weight of it in a known

volume of water. Then the given sodium hydroxide solution is standardized by titrating a known

volume of this standard oxalic acid solution against sodium hydroxide solution taken in the

burette. Since sodium hydroxide is a strong base and oxalic acid is a weak acid phenolphthalein

is used as the indicator. The above titration is repeated with the given oxalic acid solution.

MATERIALS AND METHODS:

a) Preparation of standard oxalic acid solution:

1. About 1.5g of pure oxalic acid crystals are weighed out accurately into a 250 mL standard flask.

2. It is dissolved in a little distilled water and then made up to the mark with distilled water and shaken

well for uniform concentration.


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b) Standardisation of sodium hydroxide solution

1. 25 mL of the prepared oxalic acid solution is pipettedd out into a clean conical flask.

2. Three drops ofphenolphthalein indicator is added to it.

3. The burette is filled with the given sodium hydroxide solution.

4. The initial burette reading is noted.

5. The solution in the conical flask is then titrated against sodium hydroxide solution in the burette

carefully in the beginning and dropwise towards the end of the titration.

6. Just appearance of a permanent pale pink color marks the end point.

7. Final burette reading is noted.

8. Titration is repeated to get the concordant values.

C)Estimation of oxalic acid in the given solution

1.The above titration is repeated with the given oxalic acid solution instead of the prepared standard

oxalic acid solution.

2. The concordant value obtained is noted.

Result

The weight ofoxalic acid crystals dissolved per litre of the solution =

Observations and CalculationsPreparation of standard oxalic acid solution

Weight of the weighing bottle + oxalic acid crystals, W1 =

Weight of the weighing bottle after transferring the crystals, W2 =

Weight of oxalic acid crystals taken, ( W1 - W2) =

Strength of the solution prepared, N1 = ( W1 - W2)*4 / 63

==

Standardisation of sodium hydroxide solution

Solution taken in the burette given NaOH solution

Solution taken in the conical flask 25 mL of prepared oxalic acid solution

Indicator used 3 drops ofphenolphthalein


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Colour change colourless to pale pink

Trial No. 1 2 3

Final Burette readingInitial Burette reading

Volume of NaOH added

Agreeing value, V1 =

Strength of the given NaOH solution, N2 = 25*N1/V1

=

=

Estimation of oxalic acid

Solution taken in the burette given NaOH solution

Solution taken in the conical flask 25 mL of given oxalic acid solution

Indicator used 3 drops ofphenolphthalein

Colour change colourless to pale pink

Trial No. 1 2 3



Volume of NaOH added


Strength of the given oxalic acid solution, N3 = V2 * N2/ 25

=

=

Weight of oxalic acid crystals present per litre of the solution = N3 * 63

=

=


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Estimation of Hardness of Water

INTRODUCTION:

Hardness in water is caused by dissolved salts mostly, calcium and mag

Aim of the Experiment

To estimate the total hardness of the given sample of water, by EDTA method.

Principle

Water soluble salts (bicarbonates, chlorides and sulphates) of calcium and magnesium (less

frequently iron and aluminium) is called hard water. Bicarbonates produce temporary hardness,

while chlorides and sulphates produce permanent hardness. Temporary and permanent hardness

together are called total hardness. Hardness is expressed in parts of CaCO 3 equivalents permillion parts (ppm) of water, or mg of CaCO3 equivalents per litre or dm

3 of water.

EDTA, is insoluble in water. Its disodium salt is readily soluble in water. EDTA combines with

all metal ions in 1:1 ratio and forms highly soluble complex ions. Thus, when EDTA is added to

hard water, Ca2+ and Mg2+ ions combine with EDTA and form highly soluble complex ions.

Removal of H+ ions produced, increases the stability of the complex ions formed. This is

achieved by maintaining pH of the solution in basic range using an appropriate buffer solution.

Eriochrome BlackT (EBT) a triprotic acid, is a suitable indicator for this titration and is known

as a metal ion indicator.

When EBT indicator is added to hard water at pH 10, it forms red complex ions with a small

fraction of Ca2+

and Mg2+

, depending upon the quantity of indicator used. When EDTA is added

to this solution, it complexes with all the free Ca2+

and Mg2+

ions present. When enough EDTA

is added to reaction mixture, colour changes from red to blue at this point.

A standard solution of CaCO3 is prepared by dissolving a known weight of it in a known volume

of water. A known volume of this is titrated against given EDTA solution using Eriochrome

BlackT indicator. A buffer solution of NH4Cl-NH4OH is added to maintain pH at 10. Thus EDTA

is standardized. Experiment is repeated with water sample. From the titre value, molarity of water sample

and hence hardness of water is calculated.

Procedure

Preparation of Standard CaCO3 solution:


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CaCO3 solid is accurately weighed into a 250ml beaker. Dilute HCl is added slowly till the solid

completely dissolves. The solution is just neutralized with NaOH solution till a white precipitate is

formed. Then the precipitate is dissolved in minimum amount of dilute HCl. The solution is made up to

the mark in 250ml standard flask with distilled water and shaken well for uniform concentration.

Standardisation of EDTA solution

25ml of prepared calcium carbonate solution is pipettedd into a clean conical flask. 2ml of

NH4Cl-NH4OH buffer solution is added followed by 3 drops of Eriochrome blackT indicator. The

solution turns to wine red colour. This is titrated against EDTA solution taken in the burette. Titration is

repeated to get concordant values.

Estimation of hardness of water

25ml of given hard water is pipettedd out into a clean conical flask. 2ml of NH4Cl-NH4OH buffer

solution is added followed by 3 drops of Eriochrome blackT indicator. The solution turns to wine redcolour. This is titrated against EDTA solution taken in the burette. The end point is indicated when the

solution turns to blue colour without reddish tinge. Titration is repeated to get agreeing values.

Result

Total hardness of the given sample of water =

Observations and CalculationsPreparation of standard Calcium Carbonate solution

Weight of the weighing bottle + CaCO3 crystals, W1 =Weight of the empty weighing bottle, W2 =

Weight of CaCO3 crystals taken, W = ( W1 - W2) =

Standardisation of EDTA solution

Solution taken in the burette given EDTA solution

Solution taken in the conical flask 25 mL of CaCO3 solution + 2 mL of buffer solution

Indicator used 3 drops of Eriochrome blackT

End point wine red to blue

Trial No. 1 2 3



Volume of EDTA added



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Estimation of hardness of water

Solution taken in the burette given EDTA solution

Solution taken in the conical flask 25 mL of hard water + 2 mL of buffer solution

Indicator used 3 drops of Eriochrome blackT

End point wine red to blue

Trial No. 1 2 3



Volume of EDTA added


1 mL of EDTA solution W/ (10 * V1) g of CaCO3

25 mL of water sample (W * V2) / (10 * V1) g of CaCO3

1 million of water contains (W * V2 * 106) / (10 * V1* 25) g of CaCO3

Or hardness of water = (W * V2 * 106) / (10 * V1* 25)

=

=


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Estimation of Chloride in Water


To estimate the chloride content of the given sample of water using silver nitrate solution and

pure potassium chloride crystals.

Principle

When AgNO3 solution is added to an aqueous solution containing both KCl and K2CrO4, curdy

white AgCl precipitates at first, since its solubility (1.5 x 10 -3 g dm-3 or 1.05 x 10-5 mol dm-3) is

less than that of Ag2CrO4 (2.5 x 10-2

g dm-3

or 7.5 x 10-5

mol dm-3

). Brick red Ag2CrO4

precipitates, after all Cl-ions precipitates as AgCl. Brick red colour due to Ag2CrO4 appears at

the end point of the titration.

A standard solution of KCl is prepared by dissolving a known weight of it in a known volume of

water. A known volume of this standard solution is titrated against AgNO3 solution using

K2CrO4 as the indicator. Thus the given AgNO3 is standardized.

The given water sample contains dissolved chloride salts and mineral acids. A known volume of

water is treated with solid CaCO3 to neutralize mineral acids and titrated against standardized

AgNO3 solution using K2CrO4 as indicator. From the titre value normality of water sample and hence

chloride content can be calculated. Chloride content in water sample is expressed in ppm of sodium

chloride.

This titration (Cl-

vs AgNO3 using K2CrO4 indicator) must be performed in neutral medium. In

acidic medium CrO42- ion change to Cr2O7

2- ion. In alkali medium AgNO3 precipitates as black

AgOH.

Procedure

Preparation of Standard solution of KCl:

The given KCl crystals, in weighing bottle is accurately weighed into a 250 mL standard flask. It

is dissolved in distilled water and made up to the mark and shaken well for uniform concentration.

Standardisation of AgNO3 solution

25ml of prepared KCl solution is pipettedd into a clean conical flask. 10 drops of 2% K2CrO4

solution is added as internal indicator. It is titrated against AgNO3 solution taken in the burette. End

point is marked by the appearance of reddish brown tinge, which persists even after brisk swirling.

Titration is repeated to get concordant values.


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Estimation of chloride in water

25ml of water sample is pipettedd out into a clean conical flask. 1g of solid CaCO3 is added

followed by 10 drops of K2CrO4 indicator. The solution is then titrated against AgNO3 solution as before.

Titration is repeated to get agreeing values.

Result

Chloride content in the given sample of water =

Observations and CalculationsPreparation of standard KCl solution

Weight of the weighing bottle + KCl crystals, W1 =


Weight of KClcrystals dissolved per 250 mL solution, W = ( W1 - W2)

=

=

Strength of KCl solution prepared, N = W*4/74.5

=

=

Standardisation of AgNO3 solution

Solution taken in the burette given AgNO3 solution

Solution taken in the conical flask 25 mL of KCl solution

Indicator used 10 drops of 2% K2CrO4 solution

End point Appearance of pale reddish brown tinge

Trial No. 1 2 3



Volume of AgNO3 added


Strength ofAgNO3 solution, N1 = 25 * N / V1


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Estimation of Chloride in water

Solution taken in the burette given AgNO3 solution

Solution taken in the conical flask 25 mL of water sample + 0.1g of solid CaCO3

Indicator used 10 drops of 2% K2CrO4 solution

End point Appearance of pale reddish brown tinge

Trial No. 1 2 3



Volume of AgNO3 added

Strength ofChloride in water, N2 = V2 * N1 / 25

=

=

Chloride content of water sample = N2 * 58.5 * 1000

=

=


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Determination of COD of Industrial Waste Water


Determine the chemical oxygen demand (COD) of the given industrial water, using K2Cr2O7

solution and pure Mohrs salt crystals.

Principle

Chemical oxygen demand is a measure of the oxygen equivalent of organic and inorganic

materials in a water sample that can be oxidized by a strong chemical oxidant. COD is expressed

as mg of oxygen required to oxidize impurities present in 1000cm3

of waste water.

The oxidizable constituents of waste water include straight chain aliphatic compounds, aromatic

hydrocarbons, straight chain alcohols, acids, pyridine and other oxidizable materials. Straight

chain compounds and acetic acid are oxidized more effectively in presence of silver ions (added

as silver sulphate) as catalyst. However the silver ions become ineffective in presence of halideions (present in waste water) owing to the precipitation of silver halide. This difficulty is

overcome by treating the waste water with mercuric sulphate before analysis for COD. Mercuric

sulphate binds the halide ions and makes them unavailable.

A known volume of the waste water sample is treated with excess of acidified dichromate. Unreacted

dichromate is treated with standard ferrous ammonium sulphate solution using ferroin as indicator. A

blank titration without the water sample is performed. COD of the water sample is determined from the

difference in the titre values.

ProcedurePreparation of StandardMohrs saltsolution:

Given mohrs salt is accurately weighed into a 250 mL standard flask. About 2 test tube

dilute H2SO4 is added to the flask. The salt is dissolved in distilled water, made up to the mark

with distilled water and shaken well.

Back titration:

25ml of waste water is pipettedd out into a 250mL round bottomed flask. One tube of 1:1

H2SO4 (containing 0.5g of silver sulphate and 0.5g of mercuric sulphate) is added with constant

shaking. 25ml of the given K2Cr2O7 solution is also pipettedd slowly into it. A few porcelain bits

are added and the contents in the flask are mixed. The flask is attached with a reflux condenser.

The mixture is refluxed for one hour. Contents of the flask are then cooled to room temperature,

transferred quantitatively into a 250 mL standard flask, made up to the mark with distilled water

and shaken well.


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25ml of this solution is pipettedd into a conical flask. 3-4 drops of ferroin indicator are added and

titrated with the standard Mohrs salt solution until colour changes from pale blue to dark red.

Titration is repeated until agreeing value is obtained.

Blank Titration

In another round bottomed flask instead of waste water, 25 mL of distilled water and all other

chemicals as mentioned above are taken. It is refluxed for the same length of time. Contents of

the flask are made up to the mark in a 250 mL standard flask. 25mL of this solution is titrated

with the same standard using Mohrs salt solution to the same end point using the same indicator.

Titration is repeated until agreeing value is obtained.

Result

COD of the given industrial waste water =

Observations and CalculationsPreparation of standardMohrs salt solution

Weight of the weighing bottle with Mohrs salt, W1 =


Weight ofMohrs salt in 250 mL solution, W = ( W1 - W2) =

Concentration ofMohrs salt solution prepared, Nm = W*4/392

Blank titration

Solution in the burette Standard Mohrs salt solution

Solution in the flask 25 mL of make-up solution containing distilled water

refluxed with K2Cr2O7 solution

Indicator used 3-4 drops of ferroin

End point Change of colour from pale blue to deep red

Trial No. 1 2 3



Volume of Mohrs solution addedAgreeing value, V1 =

Back Titration

Solution taken in the burette Standard Mohrs solution

Solution taken in the conical flask 25 mL of made up solution containing waste-water

refluxed with K2Cr2O7 solution

Indicator used 3 -4 drops of ferroin


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End point Change of colour from pale blue to deep red

Trial No. 1 2 3


Initial Burette readingVolume ofMohrs solution added


Amount of oxygen consumed by 1000 mL of waste water

= 10*(V1-V2)*Nm*8*1000/25

or, COD of given waste water

=

=

Estimation of CaO in Lime


Estimation of percentage of CaO in lime, using iodine and standard Na 2S2O3 solution.


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Principle

A known weight of the lime is treated with boiling water and then with a known excess of iodine

solution. The excess of iodine left over is determined by titrating against standard Na2S2O3

solution.

A blank titre value is determined for the same volume of iodine solution added to the water by

titrating with Na2S2O3 solution. The difference in the two titre values is equivalent to volume of

iodine consumed by the lime present in the sample. The percentage of CaO present in the sample

is then calculated.

Procedure

Back Titration

About 0.2g of the given lime sample is accurately weighed into a glass stoppered conical

flask. About 4 test tubes of boiling water is added and the solution is shaken for about 30minutes and cooled. Then 25mL of 1.0 N iodine solution is added from burette and the solution

is swirled well for about 10 minutes for the reaction to complete. (Any insoluble silica present is

easily distinguished from the milky appearing lime). Then it is made up to the mark with distilled

water in 250 mL standard flask and shaken well.

25mL of this solution is titrated against standard 0.1N Na2S2O3 taken in burette, using 1 mL of

starch as the indicator, which is added towards the end of titration. Titration is repeated to get

concordant values.

Blank Titration

A blank titration is carried out by making up 25 mL of 1.0 N iodine solution (added from

burette) with distilled water in another 250mL standard flask. 25 mL of this solution is titrated

against standard Na2S2O3 solution using 1mL starch as the indicator towards the end of the

titration. Titration is repeated to get agreeing value.

Result

Percentage of calcium oxide in the given sample of lime =

Observations and Calculations

Back titration

Solution taken in the burette Standard Na2S2O3 solution

Solution taken in the conical flask Excess of iodine solution

Indicator used 1 mL of starch (to be added near the end point)


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End point Discharge of blue colour

Trial No. 1 2 3



Volume of Na2S2O3 added


Blank titration

Solution taken in the burette Standard Na2S2O3 solution

Solution taken in the conical flask 25 mL of iodine solution

Indicator used 1 mL of starch (to be added near the end point)


Trial No. 1 2 3



Volume of Na2S2O3 added


Weight of weighing bottle + lime sample, W1 =

Weight of empty weighing bottle, W2 =

Weight of lime sample taken, W= (W1W2) =

=

Volume of Na2S2O3 solution required for the experimental titration, V2 =

Volume of Na2S2O3 solution required for the blank titration, V1 =

Percentage of CaO in the lime sample = [28.04 * 10 * (V1V2) * 100 * 0.1] / [1000 * W]

=

=

Analysis of Haematite Ore Volumetrically

Aim of the Experiment:

To determine the percentage of iron gravimetrically and volumetrically.


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Principle:

Iron Volumetrically by K2Cr2O7 method

Haematite is an important ore of iron containing mainly ferric oxide, Fe 2O3 and a small amountof silica, SiO2. A known weight of the ore sample is treated with concentrated HCl. Fe2O3

dissolves and SiO2 remains. The insoluble portion is filtered and the filtrate along with washings

is made up to known volume. This filtrate solution is used for estimating iron gravimetrically and

volumetrically.

A known volume of the ore solution from the 250 mL standard flask is taken and Fe3+

is reduced

to Fe2+ using stannous solution in hot condition in the presence of concentrated HCl. The ferrous

ions in the resulting solution is titrated against standard K2Cr2O7 using diphenyl amine indicator.

Procedure:

Preparation of ore solution:

About 1.5g of the well powdered ore is accurately weighed into a 250 ml beaker. About 3

test tube of concentrated HCl is added. The beaker is covered with a watch glass and gently

heated until the ore dissolves completely, leaving behind white silica residue.

About 2 test tube of distilled water is added and further boiled for 5 mins. Then cooled to

room temperature and filtered into a 250 ml standard flask. (5% dilute HCl is used as the washliquid). The filtrate and the washings are collected in the 250 mL standard flask are made up to

the mark with distilled water and shaken well.

Preparation of Mohrs salt solution

About 10g of Mohrs salt is accurately weighed into a 250ml standard flask. The salt is

dissolved in a little dilute H2SO4.The solution is made up to the mark with distilled water and

shaken well.

Standardisation of K2Cr2O7solution

25ml of Mohrs salt solution is pipetted into a conical flask. About 8 drops of 1%

diphenylamine solution is added as the indicator. Then about 25ml of sulphuric acid

phosphoric acid mixture is added. The solution is titrated slowly with constant stirring against


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K2Cr2O7 solution in the burette, till the solution becomes bluish violet and remains permanent.

This makes the end point. The titration repeated to get agreeing values.

Estimation of Fe2O3

25ml of haematite ore solution is pipettedd out into a conical flask. test tube ofconcentrated HCl is added and heated just to boiling. To the hot solution, SnCl2 solution is added

dropwise from the burette until the yellow colour of the solution just discharges. One or two

drops are added in excess. The hot solution is rapidly cooled under the tap to room temperature

with protection from the air. 5ml of saturated mercuric chloride solution is rapidly added in one

portion and with thorough mixing. A light silky white precipitate should be obtained. If a black

precipitate is obtained or if no precipitate is obtained it shoud be rejected and the process must be

repeated with another 25ml portion of the FeCl3 solution. About 8 drops of 1% diphenylamine

solution is added as the indicator. Then about 25ml of sulphuric acid phosphoric acid mixtureis

added and titrated against K2Cr2O7 solution in the burette with constant stirring till the solution

turns to intense purple or bluish-violet colour. This marks the end point. Titration is repeated toget agreeing values.

Result

Percent of iron oxide, Fe2O3 in the given sample of haematite ore =


Preparation of standard Mohrs salt solution

Weight of the weighing bottle with Mohrs salt, W1 =Weight of the empty weighing bottle, W2 =

Weight of Mohrs salt in 250 mL solution, W3= ( W1 - W2) =

Concentration of Mohrs salt solution prepared, N1 = W3*4 / 392

=

=

Standardisation of K2Cr2O7solution

Solution in the burette given K2Cr2O7 solution

Solution in the flask 25mL of Mohrs salt solution + 25 mL of acid mixture

Indicator used 8 drops of 1% diphenylamine


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End point Appearance of bluish violet colour

Trial No. 1 2 3



Volume of K2Cr2O7 solution

added


Strength of K2Cr2O7 solution, N2 = 25 * N1 / V1

=

=

Weight of the weighing bottle + haematite ore, W4 =


Weight of haematite ore taken, W= ( W4W5) =

Estimation of FeCl3

Solution taken in the burette given K2Cr2O7 solution

Solution in the flask 25 mL of FeCl3 solution + 1/3 test tube of concentrated HCl

(heated to boiling) + SnCl2 (cooled) + 5 mL of saturated

HgCl2 + 25 mL of acid mixture

Indicator used 8 drops of 1% diphenylamine

End point Appearance of bluish violet colour

Trial No. 1 2 3



Volume of K2Cr2O7 solutionadded


Strength of FeCl3 solution, N3 = V2 * N2 / 25

Percentage of FeCl3 in the sample = N3 * 55.85 * 100 / 4 * W

Analysis of Haematite Ore Gravimetrically

Principle:

Iron gravimetrically as Fe2O3


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A known volume of the filtrate solution is boiled with concentrated HNO 3 to oxidize Fe2+

to Fe3+

and then by adding a slight excess of aqueous ammonia solution. Fe3+

is quantitatively

precipitated as Fe(OH)3. The reddish brown precipitate is separated by filtration, dried and

ignited in a silica crucible and weighed as Fe2O3.

Procedure:

25 ml of the ore solution is pipettedd out into a 400 mL beaker. 1-2 mL of concentrated

HNO3 is added and boiled gently for 5 minutes. Then the solution is diluted with about 100 mL

water and heated to boiling. Now 1:1 ammonia solution is added in a thin stream with constant

stirring until a slight excess is present. It is then boiled gently for one minute and the precipitate

is allowed to settle. Then it is filtered using Whatmann No. 41 filter paper. The precipitate is

quantitatively transferred into the filter paper and washed several times with hot 1% NH 4NO3

solution. The filter paper with the residue is transferred into a weighed silica crucible. It is dried,

ignited, cooled and weighed to a constant weight as Fe2O3.

Result

Percent of iron oxide, Fe2O3 in the given sample of haematite ore =


Weight of haematite ore taken, W =

Weight of Fe2O3 residue obtained from 25mL of the filtrate solution, W6 =

Percentage of Fe2O3 in the sample = 10 * 100 * W6 / W

=

=

Analysis of Brass Volumetrically


To estimate copper gravimetrically and volumetrically in the given sample of brass alloy.


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Principle:

A known weight of brass is dissolved in moderately concentrated nitric acid. The solution is

made up to a known volume in a standard flask.

A known volume of the brass solution is treated with excess of KI after destroying the mineralacid. An equivalent amount of iodine liberated is titrated against standard Na2S2O3 solution using

starch as the indicator.

Given thiosulphate solution is standardized using K2Cr2O7 solution using starch solution as

indicator near the end point.

Procedure:

Preparation of brass solution:

About 1g of the brass is accurately weighed into a 250mL beaker. About test tube of

concentrated HNO3 is added. The beaker is covered with a watch glass. When the vigorous

reaction is over, 2 test tubes of distilled water and 1g of urea are added and boiled for 5 minutes.It is then cooled and transferred quantitatively into a 250ml standard flask and made up to the

mark with distilled water and shaken well.

Standardisation of Na2S2O3 solution

About 1.2g of K2Cr2O7 crystals are accurately weighed into a 250ml standard flask,

dissolve and make up to the mark with distilled water and shaken well. 25ml of this solution is

pipettedd out into a conical flask. 1/3 test tube of concentrated HCl and 10ml of 10% KI solution

are added and the liberated iodine is titrated against the given Na2S2O3 solution in the burette

using 1mL of starch solution as the indicator towards the end point. Titration is repeated.

Estimation of copper in brass

25ml of the brass solution is pipettedd out into a conical flask. Na2CO3 solution is added

drop by drop until a bluish white precipitate is obtained. The precipitate is just dissolved by

adding dilute acetic acid and 5 ml is added in excess. Then 10ml of 10% KI solution is added

and the liberated iodine is titrated against the given Na2S2O3 solution in the burette using 1ml

starch solution as the indicator towards the end point. Titration is repeated.

Result

Percentage of volumetrically estimated copper in brass =


Volumetric Estimation

Preparation of K2Cr2O7 solution

Weight of the weighing bottle + K2Cr2O7 crystals, W1 =


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Weight of K2Cr2O7 crystals taken, W3= ( W1 - W2) =

Strength of K2Cr2O7 solution prepared, N1 = W3*4 / 49

Weight of the weighing bottle + brass, W4 =Weight of the empty weighing bottle, W5 =

Weight of Brass sample taken, W = ( W4W5) =

Standardisation of Na2S2O3 solution

Solution in the burette given Na2S2O3 solution

Solution in the flask 25 mL of K2Cr2O7 solution + 1/3 test tube of

concentrated HCl (heated to boiling) + 10 mL of 10 % KI

Indicator used 1 mL of starch solutionEnd point Discharge of blue colour

Trial No. 1 2 3



Volume of Na2S2O3 solution

added


Strength of Na2S2O3 solution, N2 = 25 * N1 / V1

Estimation of copper in brass

Solution taken in the burette given Na2S2O3 solution

Solution in the flask 25 mL of brass solution + Na2CO3 +

acetic acid + 10mL of 10% KI

Indicator used 1 mL of starch solution


Trial No. 1 2 3



Volume of Na2S2O3 solutionadded


Strength of brass solution, N3 = V2 * N2 / 25

Percentage of copper = N3 * 63.7 * 100 / 4 * W


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Analysis of Brass Gravimetrically

Principle:

A known volume of the brass solution is taken and copper is precipitated as cuprous thiocyanate

by adding NH4CNS solution in presence of a reducing agent like sulphurous acid. The precipitate

is collected by filtration and weighed as such after drying.


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Procedure:

50 ml of the brass solution is pipettedd out into a 250 mL beaker. Na2CO3 solution is

added dropwise until a bluish white precipitate is obtained. The precipitate is just dissolved by

adding dilute HCl. 2 test tubes of saturated sulphurous acid solution and 2 test tubes of distilled

water are added and heated nearly to boiling. To the hot liquid, 1 test tube of 10% ammoniumthiocyanate solution is slowly added with stirring until it is present in slight excess. The solution

is then digested for 30 minutes over steam bath and allowed to stand for 15 minutes. It is then

filtered through a weighed sintered glass crucible. The precipitate is washed several times with

very dilute NH4CNS solution and finally with little of alcohol. It is then dried at 110 120 C for

30 minutes and weighted.

Result

Percentage of gravimetrically estimated copper in brass =


Weight of brass sample taken, W =

Weight of CuCNS precipitate obtained from 50 mL of brass solution, x =

Percentage of copper in brass = (5 * x* 63.57 * 100) / (121 * W)

Analysis of Stainless Steel


Determine the percentage of nickel gravimetrically in stainless steel.

Principle


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A known weight of stainless steel is dissolved in concentrated HCl and nitric acid to ensurecomplete oxidation of carbon to CO2 and iron to ferric state. Crystallized salts are dissolved by

boiling with water. Fe3+ and Cr3+ also interfere. They are rendered inactive by the addition of a

soluble tartarate or citrate with which these form complex ions. From this solution, nickel is

estimated gravimetrically as nickel dimethyl glyoxime.

Procedure:

About 1g of the given stainless steel is accurately weighed into a 250 mL beaker. About

1 test tube of concentrated HCl is added and heated gently until the alloy dissolves. About

test tube of concentrated HNO3 is then added carefully and slowly. It is then boiled until the

evolution of reddish brown NO2 gas stops. 2 test tubes of distilled water is then added and heated

until the crystallized salts dissolve. It is then cooled and made up to mark in a 100 mL standard

flask. Shake well for uniform concentration.

25ml of this solution is pipetted out into a 400 mL beaker and 2 test tubes of distilled water is

added. Then 20 mL of 25% tartaric acid is added followed by 1:1 ammonia solution until thesolution is basic to litmus paper. Dilute HCl is added until the solution is acidic to litmus paper.

It is warmed to 60 80 C. Then 1 test tube of 1% dimethyl glyoxime in alcohol is added

slowly with constant stirring. NH4OH solution is then added until present in slight excess

digested for 30 minutes in a water bath. It is then filtered through a weighed sintered glass

crucible. The precipitate is washed thoroughly with hot water. Then it is dried at 110120C for

one hour, cooled in a desiccator and weighed.

Result

Percentage of nickel in the given stainless steel =


Weight of the weighing bottle + stainless steel, W1 =


Weight of stainless steel taken, W3= ( W1 - W2) =

Weight of nickel dimethyl glyoxime obtained, W4 =

Percentage of nickel in stainless steel = (4 * W4 * 58.7 * 100) / (288.7 * W3)

inorganic chemistry practical protocal

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