gudlavalleru engineering collegebsh.gecgudlavalleru.ac.in/pdf/manuals/chemistrylabmanual.pdf ·...

P.Haritha, Asst. Professor, Dept. of BS&H; Gudlavalleru Engg. College 1

GUDLAVALLERU ENGINEERING COLLEGE (An Autonomous Institute Affiliated to JNTUK, Kakinada)

Seshadri Rao Knowledge Village, Gudlavalleru – 521 356. Krishna District, Andhra Pradesh

I B.TECH. R14 REGULATION


List of Experiments

Introduction to Chemistry Lab (the teachers are expected to teach fundamentals like

Primary, Secondary Standard Solutions , Normality, Molarity, Molality etc. and

laboratory ware used, error ,accuracy, precision, Theory of indicators, use of

volumetric titrations.

1) Practice experiment-Determination of the amount of HCl using standard Na2CO3.

2) Determination of alkalinity of water sample.

3) Determination of acidity of water sample.

4) Determination of Ferrous iron by permanganometric method.

5) Determination of Ferric Iron using standard K2Cr2O7 solution.

6) Determination of Total hardness of the water sample by EDTA method.

7) pH metric titrations - Determination of concentration of HCl using glass electrode.

8) Determination of pH of the water sample by using pH meter.

9) Determination of conductivity of the water sample by using conductivity meter.

10) Conductometric titrations between strong acid and strong base

11) Determination of turbidity of the water sample by using turbidity meter.

12) Estimation of total dissolved salts in water sample.

13) Preparation of Phenol - Formaldehyde resin.


Ex.No:1 Date:

INTRODUCTION TO CHEMISTRY LAB

Analytical chemistry is mainly divided into two.

1. Qualitative analysis 2. Quantitative analysis

1. Qualitative analysis: It deals with the detection and identification of the constituents

present in the given substance whether an element, compound or a mixture.

2. Quantitative analysis: It deals with the determination of the respective amounts of

any constituents of chemical substances.

Quantitative chemical analysis is further divided into two types.

i) Volumetric analysis ii) Gravimetrical analysis .

i) Volumetric Analysis: In volumetric analysis, the substance to be estimated is made

to react in solution, with another solution of a substance of a known strength.

ii) Gravimetrical Analysis: In Gravimetric analysis, the estimation of a substance is

carried out by the process of weighing. In this analysis the component to be estimated

is converted into an insoluble compound of known chemical composition, and this

compound is separated, purified, dried and weighed.

Terms used in Volumetric Analysis:

TITRATION: It is a process of one solution from the burette to another in the conical flask,

in order to complete the chemical reaction. Out of the two solutions, one must be standard.

END POINT (OR) EQUIVALENT POINT: The point at which the colour change of the

indicator is visible to the eye is called the End Point.

TITRANT: The reagent from the burette is called as Titrant.

TITRATE: The substance being titrated is termed as Titrate.

STANDARD SOLUTION: A standard solution is one, whose concentration or strength is

known, that is the amount of the substance dissolved in one liter (1000ml) of the solution.

If a reagent is available in pure state, its standard solution is prepared by dissolving an

accurately weighed amount of it in water and making the solution to a known volume by

dilution. It is called the Direct Method. All such substances whose standard solutions are

made by the direct method are termed Standard Substances or ‘primary standards’.

Oxalic acid crystals, succinic acid, anhydrous sodium carbonate, potassium dichromate,


ferrous ammonium sulphate (Mohr’s salt), silver nitrate, sodium chloride, potassium chloride,

etc. belong to this category.

A standard substance is required to fulfill the following conditions:

1. It must be obtainable in a highly purified state

2. It must be stable in air.

3. It must be readily soluble in water.

4. It should possess a large equivalent weight so as to minimize the weighing error.

5. The error in determining the end point must be negligible.

The direct method is not possible in case of those substances which, in general, fall short of

the above conditions eg. Alkali hydroxides, inorganic acids (like HCl, H2SO4, HNO3, etc.).

These are called the Secondary Standards. Their standard solutions are prepared by the so

called Indirect method. An approximately desired weight of the substance is dissolved in

water and the solution is made up to a known volume. The exact strength of the solution is

then determined by Standardisation, i.e., titration with a suitable standard reagent.

The concentration or strength of a solution can be represented in different ways.

Eg. Molarity, Normality, Molality.

Molarity (M): The molarity of a solution is defined as the no. of moles of the solute present

in one litre of solution.

SolutionofslitreofNo

molesofNoMMolarity.

.)(

mlinsolutionofVolsoluteofwtmoleculargram

SoluteofWeight.

1000.

in volumetric analysis, Molarity expression is 2

2

1

11 2nMV

nMV

where V1, M1 and n1 represent the volume, molarity and no. of moles of the first solution.

sssV2, M2 and n2 represent volume, molarity and no. of moles of the second solution.

Normality (N): The Normality of a solution is defined as the no. of gram equivalents of the solute

present in one litre of solution.

Normality (N) solutionoflitersofNo

SoluteofsequivalentgmofNo.

...

mlinsolutionofVolsoluteofwtequivgram

SoluteofWeight.

1000..


in volumetric analysis, Normality expression is V1N1 = V2N2

Where V1 and N1 represent the volume and normality of the first solution.

V2 and N2 represent volume and normality of the second solution.

Molality (m): Molality is the number of moles of the solute present per 1000 grams (1 kg) of the

solvent and is denoted by ‘m’.

Laboratory Ware used:

1) Pipette: It is used to transport a measured volume of liquid.

2) Burette: It is a vertical cylindrical piece of laboratory glassware with a volumentric graudation

on its full length and a precision tap, or stopcock, on the bottom. It is used to dispense known

amounts of liquid in experiments.

3) Conical Flask: It is a widely used type of laboratory flask which features a conical base and

cylindrical neck.

4) Volumetric Flask: It is a piece of laboratory glassware used in analytical chemistry for the

preparation of solutions. It is made of glass or plastic and consists of a flat-bottomed bulb with

a long neck, usually fitted with a stopper.

General Precautions:

1) Do not touch the bulb of pipette.

2) Do not blow the last drop of pipette.

3) Use pointing finger to control the flow from pipette.

4) The tip of pipette should immerse completely into solution, while taking the sample with

pipette.

5) Rinse the pipette and burette with the solution that you going to take in it.

6) Do no wash or rinse the conical flask with any solution except with distilled water.

7) Use glazed tile.

8) Take reading without any parallax errors.

9) For colorless solutions lower meniscus is considered for reading.

10) For coloured solutions upper meniscus is considered for reading.

Error: Error of a measurement is the difference between the measured value and the true value.

Accuracy: It is the closeness of a measured value to the true or accepted value.

Precision: Precision describes the reproducibility of measurements, i.e., the closeness of

results that have been obtained in exactly the same way.

Titrimetric/Volumetric Analysis & Type of Reactions:

1) Neutralisation reactions (or) Acidimetry and Alkalimetry:


a) Titration of free bases/salts of weak acids by hydrolysis with a standard acid

(Acidimetry).

b) Titration of free acids/hydrolysis of salts of weak bases with a standard base

(Alkalimetry).

2) Complex Formation Reactions: The combination of ions to form a soluble, slightly

dissociated ion/compound.

EDTA is the most important reagent used in titrimetric analysis of complexometry.

3) Precipitation Reactions: The combination of ions to form a simple precipitate.

Ag+ with a solution of a chloride.

No change in Oxidation state occurs.

4) Oxidation-Reduction Reactions: All reactions involving change of Oxidation number

or transfer of electrons among the reacting substances. The standard solutions are either

Oxidizing or reducing agents.

Indicator: Indicator is a substance which indicate by a sharp change in colour at the end

point in a titration.

Working(Theory) of Indicators: Indicators are of various types. These are used in

1. Acid – Base titrations 2. Complexometric titrations 3. Precipitation titrations 4. Redox titrations.

But there are some titrations in which one of the reactants itself acts as indicator. There is no need of adding indicator from outside. Such titrations are called self indicator titrations. Working of acid – base indicator: There are many acid – base indicators.

Example Colour change pH range Methyl orange Red – Orange 3.1 – 4.4

Phenolphthalein Colourless – Pink 8.3 – 10.0 Methyl red Red – Yellow 4.4 – 6.0

Litmus Red – Blue 5.0 – 8.0 Every indicator shows colour change at its respective pHrange mentioned above. For instance methyl orange shows red colour up to pH – 3.1 and orange colour beyond pH – 4.4. Methyl orange exists in two structures, one is benzenoid (yellow in basic medium) and another one is quinonoid (red in acidic medium). These two structures are in equilibrium with each other.


CH3 N N = N SO3– Yellow (Basic solution) CH3 Benzenoid

CH3 N N –– HN SO3H Red (Acid solution) CH3 Quinonoid Below pH – 3.1, it exists in quinonoid structure and shows red colour and above pH – 4.4, it exists in benzenoid structure with orange colour or yellow colour. Selection of an acid – base indicator depends mainly on two factors.

1. pH range of the indicator. 2. Change in pH of the solution at the end point.

An indicator shows colour change when the above two factors match with each other. Working of indicator in Complexometric titration: In analytical chemistry Complexometric indicators are used in Complexometric titration to indicate the exact moment when all the metal ions in the solution are sequestered by a chelating agent (most usually EDTA). Such indicators are also called metallochromic indicators. To carry out metal cation titrations using EDTA, it is almost always necessary to use a Complexometric indicator to determine when the end point has been reached. Common indicators are organic dyes such as Fast Sulphon Black, Eriochrome Black T, Eriochrome Red B or Murexide. These dyes bind to the metal cations in solution to form colored complexes. However, since EDTA binds to metal cations much more strongly than does the dye used as an indicator, the EDTA will displace the dye from the metal cations as it is added to the solution of analyte. A colour change in the solution being titrated indicates that all of the dye has been displaced from the metal cations in solution, and that the endpoint has been reached. Thus, the free indicator (rather than the metal complex) serves as the endpoint indicator. EBT + M2+ (Ca2+, Mg2+) (EBT) M2+ Wine red (EBT) M2+ + EDTA (EDTA) M2+ + EBT Wine red Colourless Blue


Working of indicator in Precipitation titration: Indicator is precipitated in readily visible form from the solution at or near the equivalence point of a titration. Ex: During the estimation of chlorides in water, potassium chromate (yellow) is used as indicator. Potassium chromate is precipitated as brick red coloured silver chromate at the end point. K2CrO4 + 2AgNO3 2KNO3 + Ag2CrO4 Yellow Brick red Working of indicator in Redox Titrations: I. Redox Indicators A. Redox titrations Using Colored Titrant 6 Fe+2 + 14 H+ + Cr2O7

-2 6 Fe+3 + 2 Cr+3 + 7 H2O Titrant No Color -----> Orange 5 Fe+2 + MnO4

- + 8H+ 5 Fe+3 + Mn+2 + 4 H2O Titrant No Color -----> Pink

B. Redox Indicators 1. Highly colored substances that may be reversibly oxidized or reduced and change

colors upon oxidation and reduction Inox + ne- + m H+ Inred

one color another color 2. Each redox indicator changes color over a certain potential range

Hence indicator must have a transition potential corresponding closely to the equivalence point potential of the titration

Inox + ne- + m H+ Inred


Expt.No:1 Date:

ESTIMATION OF HYDROCHLORIC ACID Aim: To determine the normality and amount of Hydrochloric acid present in given 1 litre solution

using 0.05N standard sodium carbonate solution.

Apparatus: 1) Burette 2) Pipette3) Conical Flask 4) Beakers 5) Burette stand 6)

Glazed tile 7) Wash bottle.

Chemicals required: Hydrochloric acid solution ,Standard Sodium carbonate solution , Methyl

orange indicator.

Principle: Hydrochloric acid reacts with Sodium carbonate solution according to the following

equation. This is known as neutralization reaction.

Na2CO3 + 2 HCl 2 NaCl + H2O + CO2

Procedure:

1. The burette is rinsed with tap water, then with distilled water finally with the given

Hydrochloric acid solution. The burette is filled with Hydrochloric acid and the

nozzle portion is also completely filled with the solution with out any air bubbles in it.

The initial reading of the burette is adjusted to ‘Zero ml.’. The burette is clamped

vertically to a burette stand.

2. A 10ml pipette is taken. It is rinsed with tap water, then with distilled water and

finally with the given sodium carbonate solution. 10ml of Sodium Carbonate is

transferred into a clean Conical flask by means of a Pipette.

3. 1 or 2 drops of methyl orange are added to the solution. The solution turns yellow in

colour.

4. The flask is placed under the burette on a glazed tile. The HCl is added slowly while

shaking the flask. The addition is continued till the color changes from yellow to

pink. It is the ‘End Point’. Just before the end point, any drops of the solution

adhering to inner walls of the flask are washed down into the flask with a few drops of

distilled water.

5. The final reading of the burette is noted. The difference between the two readings

gives the volumes of Hydrochloric acid rundown. The contents of flask are thrown

away.


6. Again 10 ml sodium carbonate solution is transferred to the flask and is titrated in a

similar way. The titrations are repeated till two consecutive reading coincide. The

readings are entered in a tabular form.

Observations & Calculations:

S.No. Volume of Na2CO3

solution in ml. ‘V2’

Burette readings Volume of HCl \

rundown ‘V1’ ml. Initial Final

V1N1 = V2N2

V1 : Volume of Hydrochloric acid

N1 : Normality of Hydrochloric acid

V2 : Volume of Sodium Carbonate solution

N2 : Normality of Sodium Carbonate solution

1

221 V

NVN

Amount of HCl present in 1 liter = Normality x Gram molecular weight of HCl.

= N1 x 36.5

Result: 1) The normality of the given HCl (N1) = N

2) Amount of HCl present in given 1 liter solution = gms


Expt.No:2

Date:

DETERMINATION OF TOTAL ALKALINITY OF WATER SAMPLE

Aim: To determine the total alkalinity of the given water sample by using 0.02N H2SO4.

Apparatus: Burette, Pipette, Conical Flask, Beakers, Wash Bottles, Burette Stand, Glazed

tile.

Chemicals Required: Standard H2SO4 solution, methyl orange indicator, phenolphthalein

indicator, water sample.

Chemical equations:

OH- + H+ H2O

CO32- + H+ HCO3

-

HCO3- + H+ H2CO3 H2O+CO2

Theory:

The alkalinity of water is due to the presence of hydroxides, bicarbonates and

carbonates, and the total alkalinity is the sum of the alkalinities caused by one or more of

these present in the water sample.

Titration to pH=8.3 with phenolphthalein or the disappearance of pink colour will

indicate phenolphthalein alkalinity.

OH- + H+ → H2O ------- 1

CO32- + H+ → HCO3

- ------- 2

Titration to pH=4.5with methyl orange or the appearance of pink colour will indicate

total alkalinity.

HCO3- + H+ → H2CO3 → H2O+CO2 ------ 3

Formula:

The amount of alkalinity interms of CaCO3 equivalents

= ppmVsampleofvolume

SOHofConcvalueTitre

s )(100050. 42

Procedure:

Part-A: Phenolphthalein Alkalinity (or) Partial Alkalinity::

1. The burette is first cleaned with tap water and then with distilled water.

2. Finally, it is rinsed with given H2SO4 solution. Then it is filled with H2SO4 Solution.

3. The H2SO4 is allowed to run out to fill the nozzle portion without any air bubble.


4. The initial reading of the burette is adjusted to zero mark and fixed vertically to the

stand.

5. First the pipette is rinsed with distilled water.

6. Pipette out 20ml of water sample into a conical flask.

7. Add few drops of Phenolphthalein indicator.

8. The color of conical flask solution is pink.

9. Titrate water sample against H2SO4 solution until the color of the solution changes

from Pink to Colourless.

10. This is the endpoint of the titration.

11. The titrations are repeated until concurrent readings are obtained.

Part-B: Total Alkalinity:


2. Finally, it is rinsed with given H2SO4 solution. Then it is filled with H2SO4 Solution.

3. The H2SO4 is allowed to run out to fill the nozzle portion without any air bubble.

4. The initial reading of the burette is adjusted to zero mark and fixed vertically to the

stand.


6. Pipette out 20ml of water sample in to a conical flask.

7. Add few drops of methyl orange indicator.

8. The color of conical flask solution is yellow.

9. Titrate water sample against H2SO4 solution until the color of the solution changes

from yellow to pink.


11. The titrations are repeated until concurrent readings are obtained.

Observations and Calculations:


Determination of Phenolphthalein Alkalinity (or) Partial Alkalinity:

S.No. Volume of water

sample (ml)Vs

Burette reading (ml) Volume of H2SO4 Solution

run down V1ml Initial Final

Phenolphthalein Alkalinity in terms of CaCO3equivalent = ppmV

V

s

10005002.01

= ………………….ppm

Determination of total alkalinity:


sample (ml) Vs

Burette reading (ml) Volume of H2SO4Solution

run down V2 ml Initial Final

Total Alkalinity in terms of CaCO3 equivalents = ppmV

V

s

10005002.02

= ………………..ppm

Result:

Phenolphthalein alkalinity of the given water sample= ------------------- ppm

Total alkalinity of the given water sample = –––––––––– ppm


Expt.No:3

Date:

DETERMINATION OF ACIDITY OF WATER SAMPLE

Aim: To determine the total acidity of the given water sample by using 0.02N NaOH

solution.

Apparatus: Burette, Pipette, Conical Flask, Beakers, Wash Bottle, Burette Stand.

Chemicals required: Standard sodium hydroxide(0.02N), Phenolphthalein indicator, Methyl

orange indicator, sodium thiosulphate, water sample.

Theory: The acidity of a solution is a measure of its capacity to neutralize bases.

Acidity is due to the presence of mineral acids like H2SO4, HCl, HNO3 and dissolved CO2 in

the form of H2CO3. These acids can be estimated by titration against standard sodium

hydroxide using methyl orange and Phenolphthalein indicators.If methyl orange indicator is

used in the titration (at pH=4.5), it gives acidity of mineral acids only. This acidity is called a

partial acidity or methyl orange acidity. If Phenolphthalein indicator is used (at pH=8.3) in

the titration, it gives acidity of all compounds(mineral acidity and CO2). This acidity is called

total acidity or Phenolphthalein acidity. The difference of these two is equal to carbonic

acid acidity. Interference due to the presence of residual Chlorine is removed by adding two

drops of Na2S2O3(hypo) solution to the water sample.

Procedure:

(A)Titration of water sample using Methyl orange indicator (Partial Acidity)


2. Finally, it is rinsed with given NaOH solution. Then it is filled with NaOH Solution.

3. The NaOH is allowed to run out to fill the nozzle portion without any air bubble.

4. The initial reading of the burette is adjusted to ‘0’ mark and fixed vertically in the

stand.



7. Add 2 drops of hypo and two drops of methyl orange indicator.

8. The color of conical flask solution becomes pink.


9. Titrate water sample against NaOH solution until the color of the solution changes

from pink to Yellow.



Determination of Partial Acidity:


sample (ml)Vs

Burette reading (ml) Volume of NaOH Solution


Partial or Methyl Orange acidity in terms of CaCO3 equivalents = )(

100050.sVsampleofvolume

NaOHofconcvalueTitre

= ppmV20

10005002.01

= ppm or mg/litre.

(B)Titration of water sample using Phenolphthalein indicator (Total Acidity)


2. Finally, it is rinsed with given NaOH solution. Then it is filled with NaOH Solution.

3. The NaOH is allowed to run out to fill the nozzle portion without any air bubble.

4. The initial reading of the burette is adjusted to ‘0’ mark and fixed vertically in the

stand.



7. Add 2 drops of hypo and two drops of Phenolphthalein indicator.

8. The color of conical flask solution is colourless.


9. Titrate water sample against NaOH solution until the color of the solution changes

from colourless to Faint pink.



Determination of Total Acidity:


sample (ml) Vs

Burette reading (ml) Volume of NaOH Solution


Total or Phenolphthalein acidity in terms of

CaCO3 equivalents = )(

100050.Vssampleofvolume

NaOHofconcvalueTitre

= ppmV20

10005002.02

= ……………. ppm or mg/litre.

(C) Acidity due to carbonic acid=Total acidity-Partial acidity.

= ppm or mg/litre

Result:

Partial acidity of water sample =

Total acidity of water sample =

Carbonic acidity of water sample =


Expt.No:4

Date:

DETERMINATION OF FERROUS IRON BY PERMANGANOMETRIC

METHOD Aim: To determine the Normality and amount of Ferrous Iron present in given 250ml

solution using Potassium Permangamate solution.

Apparatus: 1) Burette 2) Pipette 3) Conical flask 4) Beakers 5) Burette stand

6) Wash bottle 7) Glazed tile 8) Measuring jar

Chemicals required: 1. Standard Oxalic acid solution , 2. KMnO4 solution. 3. Dilute

H2SO4. 4. Mohr’s salt solution.

Theory: Standardisation of KMnO4 solution is done by using standard oxalic acid solution

and then estimation of mohr’s salt is done by using KMnO4 solution.

The chemical reaction is:

2 KMnO4 + 3 H2SO4 + 5 (COOH)2 K2SO4 + 2 MnSO4 + 8 H2O + 10 CO2

This is a redox titration. No external indicator is required. KMnO4 acts as self indicator.

Procedure:

1.Preparation of standard oxalic acid solution:

Accurately weigh the oxalic acid solid ( to prepare 0.05N) and transfer it to 250ml

volumetric flask. Dissolve the sample in little distilled water. Make up the solution with

distilled water till the lower meniscus of the solution touch the mark on the stem of the

standard flask, shake the flask well for uniform concentration.

Wt of oxalic acid transferred to 250ml volumetric flask= g

Normality(N1) of oxalic acid=mlinVacidoxalicofequg

acidoxalicofWt 1000..

.

250

100063

wt

= ……………N

2. Standardisation of KMnO4 solution:

Principle: KMnO4 is an oxidizing agent and it oxidizes oxalic acid in the presence of

dil.H2SO4 to CO2 and H2O. The chemical reaction involved in this redox titration.

2 KMnO4 + 3 H2SO4 + 5 (COOH)2 K2SO4 + 2 MnSO4 + 8 H2O + 10 CO2


1) Rinse and fill the burette with the given KMnO4 solution.

2) Rinse the pipette with the standard Oxalic acid solution and pipette out 10ml of it in a

washed conical flask.

3) Add 5ml of dilute Sulphuric acid to the solution with the help of measuring jar.

4) Heat the flask to 60 – 70oC and titrate the solution against KMnO4in the burette till a

permanent pale pink color is obtained.

5) Note the burette reading and repeat the titration till two consecutive readings coincide.


S.No. Volume of

Oxalic acid V1 (ml)

Burette readings Volume of

KMnO4 rundown

V2 (ml)

Initial

(ml)

Final

(ml)

V1N1 = V2N2

Where V1 = Volume of Oxalic acid=10ml

N1 = Normality of Oxalic acid= N

V2 = Volume of KMnO4 solution= …… ml

N2 = Normality of KMnO4 solution=?

2

112 V

NVN

= ...........N

Estimation of Mohr’s salt solution.

Principle: Potassium Permanganate oxidizes Ferrous Sulphate present in the Mohr’s salt in

the presence of dil.H2SO4 to Ferric Sulahate according to the following chemical reaction.

2KMnO4 + 8H2SO4 + 10FeSO4 K2SO4 + 2 MnSO4 + 5 Fe2(SO4)3 + 8H2O

1. Rinse and fill the burette with the standard KMnO4 solution till the upper meniscus of

the KMnO4 solution touches the zero mark on the burette.


2. Rinse the pipette with the given Mohr’s salt solution and pipette out 10ml of it in a

washed conical flask.

3. Add 5ml of dilute Sulphuric acid to the solution in a conical flask.

4. Titrate the Mohr’s salt solution against KMnO4 solution till the solution attains

permanent pale pink colour..

5. Note the final reading of the burette.

6. Repeat the titration till two consecutive readings coincide. Tabulate the readings.


S.No. Volume of

Mohr’salt V3 (ml)

Burette readings Volume of

KMnO4 rundown

V2 (ml)

Initial

(ml)

Final

(ml)

V3N3 = V2N2

Where V3 = Volume of Mohr’s salt solution=10ml

N3 = Normality of Mohr’s salt solution=?

V2 = Volume of KMnO4 solution=….. ml

N2 = Normality of KMnO4 solution= ….. N

3

223 V

NVN

Amount of Ferrous Iron present in the given 250 ml solution =

= N3 x g.equ. wt. of Iron x 250 / 1000

=…………. X 56 X 250 / 1000

Result: 1) The normality of the given Mohr’s salt solution (N3) = …………….N

2) Amount of Ferrous Iron present in given 250 ml solution = ……………g


Expt. No:5

Date:

ESTIMATION OF FERRIC IRON USING STANDARD K2Cr2O7

SOLUTION

Aim: To estimate the amount of ferric iron present in the given 100ml solution using

potassium dichromate as an intermediate solution.

Apparatus: weighing bottle, volumetric flask, funnel, burette, pipette, conical flask and

wash bottles.

Chemicals required: Ferric chloride salt, potassium dichromate, stannous chloride,

mercuric chloride, concentrated hydrochloric acid, concentrated sulphuric acid,

orthophosphoric acid and diphenylamine indicator

Determination of ferric iron:

Chemical reaction:

퐶푟 푂 + 6퐹푒 + 14퐻 → 6퐹푒 + 2퐶푟 + 7퐻 푂

Theory: Ferric ion is first reduced to ferrous ion with the help of stannous chloride.

2FeCl3 + SnCl2 ⟶ 2FeCl2 + SnCl4

The excess unreacted Stannous Chloride is oxidized by the addition of few drops of mercuric

chloride solution.

SnCl4 + 2HgCl2 ⟶ SnCl4 + Hg2Cl2

The precipitated mercurous chloride settles down. Now the solution is titrated with standard

potassium dichromate solution and concentration of the ferric ion is estimated.

Estimation of Fe3+ present in given solution: Pipette out 10ml of ferric iron solution into a

clean conical flask. Add 5ml of concentrated hydrochloric acid and heat solution to boiling.

Now add stannous chloride solution drop wise until the solution becomes colorless. Cool the

solution under tap water. Now add, drop by drop, mercuric chloride solution until a white

silky precipitate is formed. Avoid adding excess of SnCl2. If excess has been added a grey

precipitate is formed, reject the solution and start the reduction process once again. Dilute the

solution to about 100ml with distilled water then add 5ml concentrated sulphuric acid and

2ml of phosphoric acid and 3 to 4 drops of diphenylamine indicator. Note down the initial

burette reading. Titrate the solution against standard potassium dichromate solution until

violet color persists. Note down the final burette reading. Repeat the titration for concurrent

readings.


Observations and calculations

Volume of K2Cr2O7 salt solution = V1 = ……ml

Normality of K2Cr2O7 salt =N1 = …….N

Volume of Ferric chloride salt solution = V2 = 20ml

Normality of Ferric chloride = N2 =?

N1V1 =N2V2

N2 =N1V1/V2

= …..N

Amount of Fe3+ present in the given 100 ml solution =

1000 Volume(ml) X Normality X wt.euivalent Gram

= = -----------g

Report: Amount of Fe3+ present in the given 100 ml solution = …….g

S.

No

Volume of Ferric

chloride salt

solution taken(ml)

Burette

readings

(ml)

Volume of

K2Cr2O7

run

down(ml) initial final

1.

2.

3.


Expt. No:6

Date:

ESTIMATION OF TOTAL HARDNESS OF WATER SAMPLE BY

EDTA METHOD Aim: To determine the total hardness of the given water sample by using EDTA solution.

Apparatus: Burette, Pipette, Conical Flask, Beakers, Wash Bottles, Burette Stand, Glazed

Tile.

Chemicals required: Standard hard water solution(0.02N),Buffer (PH=10), EBT indicator,

EDTA solution

Introduction: Water is classified as soft water and hard water.

Soft water gives lather readily with soap. But hard water does not give lather readily with

soap due to dissolved salts like CaCl2, CaSO4, MgCl2, MgSO4, Ca(HCO3)2 and Mg(HCO3)2.

The Ca2+ or Mg2+ can be estimated by titrating with EDTA solution using Eriochrome

black-T indicator at range of pH=10.

Theory: Disodium salt of EDTA solution react with hard water to form complex with

EDTA solution. When the titration is carried out in the presence of Eriochrome black-T,

buffered to pH value 10, combines with Ca2+/Mg2+ to form a complex of wine red colour.

Chemical equations:

[Ca2+, Mg2+] + Eriochrome black-T [Ca2+,Mg2+ Eriochromeblack-T]

Hard water (Unstable and wine red color)

This when titrated with EDTA solution, quickly unstable complex dissociates and forms a

stable complex and complex free Eriochrome black-T indicator available in sky blue colour.

[Ca2+,Mg2+ Eriochrome blck-T] + EDTA [ Ca2+,Mg2+ EDTA] + EBT

wine red (Stable and colorless) (blue)

Change of wine red colour to sky blue indicates the end point of titration. From the value of

EDTA solution consumed, the total hardness of water can be calculated in terms of CaCO3

equivalents

1.Standardisation of EDTA solution:

Procedure:

1. Rinse and fill the burette with the given EDTA solution.

2. Transfer 20ml of standard hard water sample into a clean conical flask.


3. Add 1ml of pH=10 buffer (NH4Cl+NH4OH buffer) solution and 1 or 2 drops of EBT

indicator. Solution turns to wine red colour.

4. Titrate with EDTA solution till wine red colour changes to sky blue. This is the end

point.

5. Repeat the titration until two consecutive readings coincide.


S.No.

Volume of

standard hard

water sample

(v1ml)

Burette reading (ml)

Volume of EDTA rundown

(V2ml) Initial Final

V1N1 = V2N2

V1= Volume of standard hard water =20ml

N1= Normality of standard hard water =0.02N

V2= Volume of given EDTA solution=……ml

N2= Normality of given EDTA solution=?

N2= 2

11

VNV

=

2. Determination of Total Hardness of the given water sample

1. Fill the burette with the EDTA solution.

2. Transfer 20ml of the given water sample into a clean conical flask.

3. Add 1ml of pH=10 buffer solution and 1 or 2 drops of EBT indicator. Solution

turns to wine red colour

4. Titrate with EDTA solution till wine red colour changes to sky blue. This is

the end point.

5. Repeat the titration until two consecutive readings coincide.



S.No.

Volume of

standard hard

water sample

(V3ml)

Burette reading (ml)

Volume of EDTA rundown

(V2ml) Initial Final

V3 N3 = V2 N2

V3= Volume of given water sample =20ml

N3= Normality of given water sample =?

V2= Volume of EDTA solution run down=…ml

N2= Normality of given EDTA solution=….N

N3= 3

22

VNV

=

Total degree of hardness of the given water

sample in terms of CaCO3 equivalents = Normality(N3)× 50 × 1000

= ……. ppm

Result:

Total degree of hardness of the given water sample is ….mg /l(or) ppm


Expt.No:7

Date:

PH-METRIC TITRATION: DETERMINATION OF CONCENTRATION OF HCl

USING GLASS ELECTRODE

Aim: To determine the concentration of unknown HCl by using 0.1N NaOH solution pH

metrically.

Apparatus: pH meter, electrode, beaker, pipette, burette.

Chemicals required: Hydrochloric acid (HCl), 0.1N sodium hydroxide (NaOH). Buffer

solutions of pH = 4 and 7.

Theory: All acids dissociate in aqueous solution to yield H+ ions. Some acids like HCl,

H2SO4, HNO3 etc. are completely ionized in aqueous medium where as CH3COOH, HCOOH

etc. ionize to a small extent only. The former is known as strong and the later as weak acid.

pH of any solution is defined as (–log [H+]) and has values between 0–14. pH < 7 indicate

acidic solution, pH > 7 indicate basic solution and pH = 7 means neutral solution.

The pH of a solution can be measured accurately with the help of a pH meter. Measurement

of pH is employed to monitor the cause of acid-base titration. The pH values of the solution

at different stages of acid–base neutralization are determined and plotted against the volume

of alkali added. On adding a base to an acid, the pH rises slowly in the initial stages as the

concentration of H+ ion decreases gradually. But, at the equivalence point, it increases rapidly

as at the equivalent point H+ ion concentration is very small. Then it flattens out after the end

point. The end point of the titration can be detected where the pH value changes most rapidly.

Procedure: Switch on the instrument and wait for 10–15 minutes so that machine gets

warmed up. Prepare the buffer solution by adding buffer tablets of pH = 4 and pH = 7 in 100

mL of water separately. Wash the electrode with distilled water and calibrate the pH meter as

per the instruction manual.


Clean the electrode with distilled water and wipe it with a tissue paper or filter paper. Take

20 mL of HCl solution in a 100 mL beaker and immerse the electrode in it. 0.1N NaOH

solution is filled in the burette . The reading shown on the scale of pH meter is pH value of

the HCl solution. Add NaOH solution drop wise from the burette (maximum 1mL at a time),

shake the solution well and note the corresponding pH values. Near the end point, the acid is

neutralized and there will a sharp increase in pH values. Further addition of even 0.01 mL of

NaOH, increase the pH value to about 9–10.

Plot a graph between pH and volume of NaOH added and find out the volume of NaOH

required (V2 mL) for complete neutralization of HCl from the graph.

V1N1=V2N2

Normality of HCl (N1) = (V2N2/20) (N)

Result: The concentration of the unknown HCl is ________(N)

Precautions: i) Electrodes must be immersed in the solution properly and sufficient time

to be allowed for the electrodes to obtain the temperature of the solution.

i) pH meter should be calibrated before the experiment.

i) Magnetic stirrer may be used or the solution be stirred mechanically from time to

time during pH metric titration.


Expt.No:8

Date:

DETERMINATION OF pH OF THE GIVEN WATER SAMPLE

Aim: To determine the pH of the given water sample using pH meter.

Apparatus: Digital pH meter, combined electrode, beakers.

Chemicals required: Standard buffer solution (pH =4; pH =7), distilled water, water

sample.

Theory: pH is defined as the negative logarithm of the H+ ion concentration.

pH = -log[H+]

where H+ is the concentration of hydrogen ion expressed in moles/litre

pH value indicates whether a solution is acidic or neutral or basic.

If pH is between 0 to 7, it indicates acidic solution.

If pH is in between 7to 14, it indicates basic solution.

If pH is exactly 7, it indicates neutral solution.

pH is determined by using pH meter which contains glass electrode which is

sensitive to H+ ions.

Preparation of Buffers:

i) Dissolve a pH =4 buffer tablet in 100ml distilled water. It indicates the standard solution of

pH =4

ii) Dissolve a pH=7 buffer tablet in 100ml distilled water. It indicates the standard solution of

pH =7

Procedure:

Calibration of pH meter:

i) Set the temperature knob to buffer solution temperature.

ii) Dip the electrode in buffer solution of pH =7.

iii) Set the display value to pH =7 by turning CAL knob.

iv) Wash the electrode with distilled water and dry with the tissue paper.

v) Dip the electrode in the buffer solution of pH=4 and set the slope button to display 4.

vi) Now the instrument is ready to measure pH .

Measurement of pH:

i) Wash the electrode with the distilled water and dry with tissue paper.


ii) Dip the electrode in the sample water to be analyzed.

iii) The reading obtained is the pH value of sample.

iv) Rinse the electrode before and after measurement of each sample.

Precautions:

a) The pH electrode should be handled gently

b) The electrode should be always dipped in distilled water when not in use

Significance:

As per Indian standards the pH of the drinking water should be between 6.5-8.5 . As per

WHO standards, the pH of the drinking water should be between 7.0-8.0 .

Report: The pH of the given water sample is


Expt.No:9

Date:

DETERMINATION OF CONDUCTIVITY OF WATER SAMPLE

Aim: To determine the conductivity of the given water sample by using conductivity meter.

Apparatus: Conductivity meter.

Reagent: Standard electrical conductivity solution.

Standard solution is prepared by dissolving 0.7456grams of KCl in 1 litre distilled

water. It will have 1413 µmhos/cm as conductance.

Theory: Conductivity is the capacity of the water to carry an electrical current.

Conductivity is proportional to the concentration of dissolved solids(ions) in the

solution.

Procedure:

Calibration of electrical conductivity meter:

(i) Adjust “cal” till the display reads 1000.

(ii) Set the “Range” switch to the proper range for the standard solution.

(iii) Change the “cond/cell constant” switch position to cell constant position.

(iv)Wash the electrode with distilled water and dry with tissue paper. Immerse the electrode

in standard solution.

(v) Adjust the cell constant present control on the rear panel till the instrument reads “1”. The

conductivity of the solution is1413µs/cm.

Measurement of conductivity:

(i) Wash the electrode with distilled water and dry with tissue paper.

(ii) Dip the electrode in 50ml sample for which conductivity is to be measured.

The reading indicated on the meter is the conductivity of the sample.

(iv) Rinse the electrode before and after measurement of conductivity of each sample.

Significance: Conductivity increases as the dissolved ionic species in water increases.

According to Indian standards, the conductivity should be inbetween 5-50mS/m.

Report: The conductivity of the given water sample is


Expt.No:10

Date:

CONDUCTOMETRIC TITRATION BETWEEN STRONG ACID AND

STRONG BASE

Aim: To determine the amount of the given HCl present in the given one litre solution by

using 0.1N NaOH solution conductometrically.

Apparatus required: Conductometer, Conductivity cell, Magnetic stirrer/ Glass stirrer,

beaker, burette and pipette.

Chemicals required: NaOH solution, HCl solution and distilled water

Principle: According to Kohlrausch’s law, the electrical conductance of a solution depends

upon the number and mobility of ions.For the titration of a strong acid like HCl against a

strong base like NaOH, before the NaOH solution is added, the acid solution has a high

conductance due to the highly mobile hydrogen ion. As alkali solution is added the hydrogen

ions are removed due to combination with the hydroxyl ions forming feebly ionized water

(H2O) molecules and their place is taken by the much slower moving Na+ ions.

퐻 퐶푙 + 푁푎 푂퐻 → 푁푎 퐶푙 + 퐻 푂(푓푒푒푏푙푦 푖표푛푖푧푒푑)

Consequently the conductance of the solution decreases and continues to decrease on adding

sodium hydroxide solution until equivalence point is reached. Any further addition of alkali

means increase of Na+ and fast moving hydroxyl ions and thus the conductance begins to

increase. If we plot the conductance measured against the volume of NaOH added, the point

of intersection will give the neutralization point.

Significance:

Titrations involving coloured liquids are difficult to be performed using ordinary indicators.

But these can be easily carried out by conductometric titration method. Moreover,

conductometric titration can be performed accurately even with dilute solutions

Procedure:

1. Wash the conductivity cell thoroughly with distilled water and calibrate the

instrument to its cell constant vaue.

2. Take 20 ml of the given HCl solution whose strength is to be determined in a 100 ml

beaker.


3. Clamp the conductivity cell in the beaker containing the acid solution.

4. Set the function switch to Cond. position.

5. Fill the NaOH solution in the burette.

6. Note down the conductance from the conductometer.

7. Keep on adding NaOH from the burette, 1mL at a time.

8. Stir the solution with magnetic stirrer or a glass rod and measure the conductance.

Graph: The values of observed conductance are plotted along Y axis against the volume of

NaOH added along the X axis.

The point of intersection gives the amount of NaOH required for neutralization of the HCl.


Observations and calculations:

N1V1 = N2V2

Volume of unknown HCl solution taken (V1) = 20ml

Normality of HCl(N1)=?

Volume of NaOH at end point (from graph) =….V2 ml

Normality of NaOH(N2)=0.1N

N1= 0.1V2/20

Amount of HCl present in given one litre solution = Normality X gram equivalent wt. of HCl

= N1X 36.5

= ………g/L

Result:

The amount of HCl present in the given one litre solution = ……. g.

S.No.

Volume of NaOH

consumed

v (ml)

Conductance (C)

1

2

3

4

5

6

7

8

9

10

11

12

13

14

15

16

17

18


Expt.No:11

Date:

DETERMINATION OF TURBIDITY OF WATER SAMPLE Aim: To determine the Turbidity of the given water sample using Turbidity meter.

Apparatus: Nephelometer.

Reagents:

Solution-1: Dissolve 1g of Hydrazine sulphate in 100ml distilled water.

Solution-2: Dissolve 10g of Hexamethylene tetramine in 100ml of distilled water.

Solution-3: Mix 5ml of solution-1 and 5ml of solution-2 in a 100ml volumetric flask. Allow

to stand for 24hrs. Dilute it to 100ml. Now the solution has a known turbidity of 400NTU.

From this stock solution, dilution is carried out to obtain lower

concentrations of 100, 40, 20 and 10 NTU.

Theory: Turbidity is a measure of cloudiness. Turbidity is the property of water because of

which it offers resistance to passage of light. It is caused by suspended solids, living or dead

micro-organisms etc.

The turbidity is measured from the amount of light scattered by the sample taking a

reference with standard turbidity suspension.

Procedure: Calibration of Nephelometer:

i) Select the appropriate range

ii) Insert the test tube( cuvette) with distilled water into cell holder and cover with light

shield.

iii) Adjust zero button to get zero an the display

iv) Take the standard solution of 100 NTU in the test tube and adjust calibration as 100

Measurement of Turbidity:

i) Rinse the test tube before and after measurement of each sample with distilled water and

dry it with a tissue paper.

ii) Take unknown sample of water in the test tube upto the mark indicated on it and place it

in the holder. The display directly gives the turbidity in NTU.

Significance: As per WHO standards, the turbidity of drinking water should be less than 5

NTU.

Report: The turbidity of the given water sample is ….NTU


Expt.No:12

Date:

ESTIMATION OF TOTAL DISSOLVED SALTS IN WATER SAMPLE Aim: To estimate the amount of total dissolved salts present in the given water sample.

Apparatus: China dish, analytical balance, measuring jar, hot air oven, water bath,

dessicator.

Principle: Water may contain dissolved salts like bicarbonates, chlorides, sulphates of

calcium, magnesium, sodium, potassium, iron etc.

The total dissolved salts in a water sample is estimated by evaporating in a

weighed dish and drying it at a constant temperature of 103°C to 105°C. The increase in

weight over that of the empty dish represents the total dissolved salts.

Procedure:

1. Take a clean and dry china dish, weigh it (W1).

2. Take 50ml of the given water sample with the help of measuring jar into the china

dish.

3. Evaporate the water to drying on a steam bath.

4. Keep the dish in hot air oven at 103°C to remove the residual water.

5. Cool the dish in a desiccator to room temperature.

6. Then weigh the dish along with the residue (W2).

Observations and calculations:

W1 = Weight of the empty china dish = ………

W2 = Weight of the dish along with the residual = ………

V = Volume of water taken = 50 ml.

Total dissolved salts = ppmV

WW 612 10

Significance:

Many dissolved salts cause displeasing taste, colour.

Estimation of total dissolved salts is useful to determine whether the water is useful for

agriculture , industrial purpose or not. In industries the use of water with high amount of salts

will lead to scaling, corrosion in boiler.

Result: The total dissolved salts present in the given water sample is


Expt.No:13

Date:

PREPARATION OF PHENOL- FORMALDEHYDE RESIN Aim: To prepare Phenol- Formaldehyde resin and calculate the yield of the product.

Apparatus: Beaker, Glass rod, Funnel

Chemicals required: Glacial acetic acid,40% Formaldehyde solution,Phenol,Concentrated

HCl solution

Theory: Phenol and formaldehyde undergo condensation polymerization to produce a three

dimensional phenol- Formaldehyde resin.

Procedure:

1. Place 5ml of glacial acetic acid and 2.5ml of 40% formaldehyde solution in a 500ml

beaker.

2. Add 2g of phenol to it.

3. Wrap a cloth loosely around the beaker. Add a few ml of conc. HCl into the mixture

carefully and heat it gently.

4. A large mass of plastic, pink in color is formed.

5. The residue is washed with water and filtered.

6. The product is dried and weighed.

Precautions:

1. Handle phenol carefully.

2. While adding conc. HCl, it is better to stay a little distance away from the beaker since

the reaction sometimes becomes vigorous.

3. The reaction mixture should be stirred continuously.

Result: The amount of Phenol Formaldehyde resin prepared= ……. grams

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