What chemical property differentiates sulfurous acid from sulfuric acid

Practical Section

A. Wet chemistry practical’s (inorganic portion)

In the following practical work , we will be confined to experiments based on wet chemistry or bench chemical analysis ; these are traditional laboratory tests which encompasses certain processes like sampling , weighing , preparing solutions and performing classical chemical tests . In ordinary wet chemistry analysis , qualitative and quantitative results can be obtained , yet it is rarely employed in areas like industry , research centers , forensic analysis etc. These institutions employ modern chemical analysis which is extension of the classical analytical techniques that has been automated and computerized to save time and increase the sensitivity selectivity of the system .

The following laboratory sessions involves certain quantitative topics like gravimetry , titrimetry and qualitative topics on identification methods based on formation of color , precipitate or evolution of specific gas .

Practical chemistry is mere laboratory work in which the theoretical knowledge of the subject is developed and the necessary skills are acquired . To undertake this work , the necessary apparatus , equipment and reagents must be available and become familiar with the students . For any given experiment , these materials must be planned and prepared so that the result can be reported in a scientific manner . The students must have knowledge of the application of these tools and get acquainted with the laboratory environment .

We can classify the laboratory tools into two parts:

A. Glassware: which are glass containers used for mixing , measuring , holding and heating laboratory reagents ; they are generally made from durable and chemically inert materials like borosilicate’s . These glass containers are of different shapes and sizes used for different purposes. Some of these laboratory glassware are stated below .

1. Beakers : Beakers are glass containers of almost cylindrical shape used for heating mixing and holding samples . They are available in different sizes .

نتيجة بحث الصور عن ‪sizes of volumetric flasks‬‏ نتيجة بحث الصور عن ‪erlenmeyer flask‬‏ نتيجة بحث الصور عن ‪sizes of volumetric flasks‬‏

Volumetric Flaska

Elenmeyer Flaska

2. Flasks : Flasks are narrow-necked and wider bottom glass containers used for measuring preparing and heating solutions . Volumetric flasks are used for preparing and diluting solutions of definite concentrations . Erlenmeyer flasks or conical flasks are used for heating or handling solutions .

نتيجة بحث الصور عن ‪boiling flask flat bottom‬‏ نتيجة بحث الصور عن ‪boiling flask flat bottom‬‏ نتيجة بحث الصور عن ‪boiling flask flat bottom‬‏

Boiling Flasks

3. Boiling flasks: flat bottom flasks used for heavy duty boiling. They are available in different forms and shapes according to their application.

نتيجة بحث الصور عن ‪test tube holder‬‏ نتيجة بحث الصور عن ‪test tube holder‬‏ نتيجة بحث الصور عن ‪test tube‬‏

Test tubes

Different boiling flasks

Test tubes holders

4. Test tubes: Small glass tubes resistant to high temperature used for testing samples especially qualitative tests .

نتيجة بحث الصور عن ‪beakers‬‏ نتيجة بحث الصور عن ‪beakers‬‏

Beakers

5. Funnels : Funnels are laboratory apparatus with a wide mouth and long pipe like opening at the end used to transfer liquid substance or finely grinded particles to avoid spillage . They are usually made of glass , plastic or ceramic .

نتيجة بحث الصور عن ‪funnels‬‏ نتيجة بحث الصور عن ‪funnels‬‏ نتيجة بحث الصور عن ‪funnels‬‏

Different funnels

6. Burettes: Burettes are long graduated glass tubes with stopcock at end for drawing accurate amount of a solution . It is extensively used in analytical chemistry, there are different sizes and structures

نتيجة بحث الصور عن ‪burette‬‏ نتيجة بحث الصور عن ‪burette‬‏

Different burettes

7. Other laboratory glassware: Pipettes are used for taking fixed volume of a liquid. They are usually made of glass or plastic materials.

نتيجة بحث الصور عن ‪pipette‬‏ نتيجة بحث الصور عن ‪pipette‬‏

Uni-Volume Pipette Graduated Pipettes

نتيجة بحث الصور عن ‪wash bottle‬‏ نتيجة بحث الصور عن ‪filtering crucible‬‏ نتيجة بحث الصور عن ‪pipette filler‬‏

Pipette fillers Crucibles wash bottles

نتيجة بحث الصور عن ‪glass rod‬‏ نتيجة بحث الصور عن ‪dessicator‬‏

Desiccator Glass rods
نتيجة بحث الصور عن ‪burette stand‬‏ نتيجة بحث الصور عن ‪burette stand‬‏ نتيجة بحث الصور عن ‪test tube rack‬‏

Burette clamp Burette stand Test tube rack

B. نتيجة بحث الصور عن ‪ph meter‬‏ نتيجة بحث الصور عن ‪hotplate‬‏ نتيجة بحث الصور عن ‪balance lab‬‏Instruments: The other feature of the laboratory equipment are the instruments. Instruments are devices that help or facilitate to undertake experimental work . meter hotplate electronic balance

نتيجة بحث الصور عن ‪shaker flask‬‏ نتيجة بحث الصور عن ‪lab oven‬‏ نتيجة بحث الصور عن ‪mortar and pestle‬‏ نتيجة بحث الصور عن ‪bunsen burners‬‏

Mortar and pestle Bunsen burners Oven Shaker

نتيجة بحث الصور عن ‪sizes of volumetric flasks‬‏Gravimetric analysis

One of the most accurate and precise methods in chemical analysis is the gravimetric analysis . In this method , the analyte is converted to insoluble form by precipitating from the solution . The precipitate is then separated , washed , dried or ignited , weighed and calculated . Perfect example can be taken in the determination of Cl- content in a given solution by precipitating as AgCl with silver nitrate solution . After filtering and drying , the precipitate is weighed and the chloride content is calculated .

There are different forms of gravimetric analysis , and the one we are concerned in this discussion is the one in which the sought for substance is almost completely precipitated from a solution so that no considerable amount is lost during filtration , washing and weighting .To achieve successful analytical result , the following points has to be fulfilled :

A. The precipitated substance must be sufficiently insoluble so that we can analytically ignore the dissolved amount .

B. The precipitate particles must be large enough to be quantitatively filtered and the impurities can easily be washed .

C. The final precipitate must be a substance of definite chemical composition after physical or chemical treatments.

· To accomplish these points and obtain almost fine precipitate particles, the following problems must be overcome too.

1. Co-precipitation, the precipitate formed may contain foreign impurities from the solution which depends on the nature of the precipitate and the conditions in which the precipitate was formed ; this contamination can be due to either the adsorption of impurities on the surface of the precipitated particles in the solution or the occlusion of foreign particles during the process of crystallization .

2. Coagulation, is the tendency of adsorption layers of the precipitate to attract ions of opposite charge forming large size particles containing solvent molecules .

3. Occlusion, this is the way in which foreign particles are trapped inside the precipitate during the crystallization . It is not easy to get rid of the occluded impurities .

4. Post-precipitation, it is the precipitation that takes place on the surface of the precipitating reagent , usually it comes from a substance that has common ion with primary precipitate , e. g CaC2O4 / Mg++ & C2O4- - .

Treatment of the Precipitate

To prepare relatively fine pure precipitate , the following favorable conditions has to be fulfilled.

A. The precipitation must be implemented in dilute solution keeping in mind the solubility of the precipitate and digestion period , to minimize the co-precipitation.

B. Add the reagents together slowly with constant stirring to create large crystals.

C. Precipitation is done in hot solution provided no counter effect in terms of solubility and the stability of the precipitate is resulted.

D. The precipitate should be washed with a suitable solvent.

Precipitation from Homogenous solution

We have already seen that to create a desirable precipitate , we must minimize the degree of supersaturation so that dilute solution of the precipitating agent is gently and slowly added to the sample solution with a constantly effective stirring . Even though these techniques maintain a low degree of supersaturation yet excess of the precipitating reagent is unavoidable so a technique known as homogenous precipitation is employed to avoid undesirable effects. In this method, the precipitating reagent is generated in situ through chemical process which takes place uniformly inside the solution. Dense precipitate is formed in this process which can easily be filtered.

There are many different anions that can be generated in this technique like hydroxyl (OH -) , phosphates (PO43-) , oxalates (C2O42-) and sulfates (SO42-). An example can be taken as the case of precipitating hydrous iron oxide or aluminum oxide. The precipitating reagent is produced by the hydrolysis of urea in lower pH .

(NH2)2C=O + H2O CO2 + NH4+ + OH –

This reaction takes place gently below the boiling point of water. Another example of homogenous precipitation is the generation of SO42- to precipitate barium (Ba) or lead (Pb) . The sulfate is generated in situ by heating Sulfidic acid solution that hydrolyzes as following :

NH2SO3H + H2O H+ + SO42- + NH4+

· Experiment (1): Gravimetric determination of Cl.

Apparatus and reagents

A. 0.1N AgNO3 (dissolve 10.79 g in dilute nitric acid and dilute to 1liter ) .

B. Conc. HNO3

C. Conc. NH3 solution

D. 3N HCl (take 252 ml of 37% HCl and dilute to liter).

E. Filtering crucibles

Procedure:

Prepare two filtering crucibles by cleaning them with conc. HNO3 followed by washing with distilled water and then with conc. NH3 followed by plenty of distilled water . Record the weight of the crucibles to three decimal places .

Take 25ml of a dilute sodium chloride solution into each of two 250ml beaker, add excess of silver nitrate solution slowly and constantly stirring . Heat almost to boiling for about ten minutes. Cool the solution and add few drops of the silver nitrate to make sure the precipitation has completed. Let the two beakers stand still for at least one hour in a dark place . Transfer the precipitate to the filtering crucibles by decanting. Wash the precipitate with few ml of 6M HNO3 followed by deionized water. Continue washing until there are no traces of AgNO3. Dry the precipitate to 110o C. for one hour. Store the crucibles in a desiccator until they are cooled to room temperature. Take the weight of the crucibles and calculate the amount of chloride in the solution. Report the % yield .

· Experiment (2): Gravimetric determination of sulfate in a solution.

Sulfate is precipitated as barium sulfate from aqueous solution by systematic addition of BaCl2 . Despite this precipitation is fast and quantitative , it has great tendency to occlude many unwanted anions such as PO43- and NO3- .

Apparatus and reagents,

A. 0.2M BaCl2.2H2O (dissolve 48.8 g in water and dilute to 1liter) .

B. 6M HCl (half-diluted 37% HCl) .

C. 2 porcelain crucibles.

Procedure : Prepare two porcelain crucibles by cleaning and igniting at high temperature. Cool the crucibles in a desiccator . Weigh the crucibles repeatedly to 0.2mg. Take two 50ml aliquots of the solution sample (SO42-) and transfer to two 400ml beakers ; add 5ml of 6M HCl followed by 100ml hot BaCl2 solution quickly with vigorous stirring. Digest the precipitate for one hour. Filter the precipitate quantitatively in an ash less filter paper. Transfer the filter papers and the precipitates into the crucibles .Ignite the contents until almost constant weight is obtained. Calculate the amount of sulfate in grams.

· Experiment (3): Determination of Barium sulfate by precipitating from homogeneous solution.

Apparatus and reagents .

A. 0.01M BaCl2.2H2O solution (A.R) [dissolve 1.22 g in distilled water and dilute to 1liter].

B. Sulfidic acid, (NH2)2SO3H (A.R)

C. Electric Hotplate

D. Filtering crucibles (Porcelain) accurately weighed to constant value.

Procedure : Take 100ml of the barium chloride solution into clean dry 250ml – beaker add 1.0g of the Sulfamic acid . Cover the beaker and heat on an electric hotplate at 980C , continue heating until turbidity begins to appear . Filter the solution in a porcelain crucible ; wash the precipitate with warm water twice . Heat the crucible with precipitate at 6500C to a constant weight and determine the weight of BaSO4 .

Titrimetric Analysis

This topic falls into a wide field that deals with quantitative determination of an analyte , it is based on practically reacting volumes in which the volume of a standard reagent required to completely consume the analyte is determined ; this procedure is known as volumetric analysis.

Volumetric procedures employ reagents whose concentrations are exactly known. These reagents are known as standard solutions; the quantity of an analyte is determined from the volume of the standard solution , this is performed by carefully adding the standard solution into the system until the reaction with the analyte is complete , then the volume of the standard reagent is measured from which , the quantitative relationship of the sought for substance is calculated . In this procedure , the end of the analyte reaction is determined by employing an indicating system to denote the equivalence point during the addition of the standard solution. This indicating technique is shown by change in color , redox potential , or self-indicating system .

Accuracy of the volumetric analysis depends on the accuracy of the preparation of the standard solution . Therefore , special attention must be made to prepare standard solution of accurate concentration . This is achieved by dissolving a carefully weighed quantity of the pure reagent and diluting to the desired volume . Highly purified substance called a primary standard solution is employed to accurately fix the real concentration of the prepared solution . This process is known as standardization .

Any substance that can be employed as primary standard solution should at least has the following properties

A. It must be Commercially available in a very high purity standard .

B. It must be least affected by heat, light, and humidity.

C. It must have high molecular weight.

Here are some of the compounds usually used as primary standard solutions :

Substance

Molecular mass

Purity available

Standardization used

Potassium dichromate, K2Cr2O7

294.22

99.9 %

Reducing agent

Potassium hydrogen phthalate(C4H5KO4)

204.23

99.9 %

Bases

Sodium carbonate, Na2CO3

105.99

99.9 %

Acids

Potassium iodate, KIO3

214.01

99.9 %

Sodium Thiosulfate

Ethylenediaminetetraacetic Acid disodium salt, 2H2O

372.25

99 %

Metals

Sodium chloride, NaCl

58.45

99.9 %

Silver Nitrate

The importance of a titration is to find out or to estimate the volume of a standard reagent chemically equivalent to a given volume of an analyte solution. The point at which the reaction stops or the analyte is theoretically consumed is called the equivalence point. The practically observed point in which the analyte is supposed to react completely is called the end point. The equivalence point is a theoretical concept while the end point is a practical concept .

In certain volumetric analysis , the end point is detected by certain chemical substances that change in color as the process approaches the equivalence point; these substances are called indicators .

· Acid–base titrations.

We have studied in details the properties and reaction of the acids and bases in chapter ( 5 ) ; but in this section we will confine ourselves to the practical aspect of the acid – base titrations . In acid – base titrations, given amount of a base solution is applied to exactly determine the chemically equivalent amount of standard acid solution , in this reaction aqueous solution of the corresponding salt is formed ; this process is called neutralization process. Different indicators are used to determine the equivalence point by changing the color of the solution . The transition from acidic color to basic color depends on the pH . This color change is not practically so abrupt but takes an interval of nearly two pH units . Color change interval varies with different indicators , so that one has to select an indicator that distinctly changes the color at the appropriate pH level .

In neutralization reaction :

H3+O + OH - 2H2O

This reaction denotes the neutralization of a strong acid and strong base . From this reaction we can say that acid is any substance that supplies hydronium ions [H+] or any substance that consumes the hydroxide ions [OH -] ; while the base is any substance that supplies hydroxide ions or consumes the hydronium ions [H3O+] . This is the core of the Arrhenius concept towards the acid – base theory.

In non-aqueous solutions , this definition is not enough and has to modified to cope with certain reactions involving Complexometric reactions . According to Arrhenius , acids furnish hydronium ions in aqueous solutions and can be stated in the following form .

HA + H2O H3+O + A - (2)

The equilibrium constant for the dissociation of the acid HA will be:

Ka =

Assuming the concentration of H2O so large with respect to other species , we can consider it . Therefore the dissociation constant of the acid will be:

Ka = (3)

We also know that in aqueous solution the dissociation constant of water ,Kw , is

Kw = [H3+O][OH -] ;

the dissociation constant of base is,

A -+ H2O HA + OH –

Kb = (4)

From equation 3 and 4 , Ka =

Substituting for the value of Kb as , Ka = , and Kb =

So Ka Kb = Kw , and pKa + pKb = Kw where pK = - log k

نتيجة بحث الصور عن ‪ph titration curve‬‏ In aqueous solutions , strong acids and strong bases ionize completely so it becomes much easier to calculate both hydrogen and hydroxyl ions from the stoichiometric reaction between the acid and the base. This quantitative process is clearly shown by using pH meter in which stepwise addition of definite volume of the acid to the base is recorded and the following pH curve is obtained . Fig.1

From these curves , the equivalence point can easily be determined . If the titration is performed without pH meter , visual indicator is used to determine the equivalence point .

Different acid-base titrations need different visual indicators and choice of the indicator is determined by the nature of the reacting acids and bases .

In strong acid-base titrations , the equivalence point fall in between pH 3.3 and 10.7 , in that case phenolphthalein will be the suitable indicator , see table ( 2 ) . The concentration of the reacting species has an effect on the inflection point. The more the concentration the sharper the inflection point .

Experimental procedure of acid – base titrations are conducted in this chapter. Before beginning detailed experiment , we have to have well prepared standardized solutions , these solutions are usually prepared from certain substance that carry the properties of primary standard solution . These substances include , Na2CO3 , H2C2O4 , H2SO4 , H3BO3 etc .

Hydrochloric acid is commonly used titrant ; although different standard solutions can be prepared , it is more convenient to prepare HCl of approximate normality and standardize with a primary standard like sodium carbonate , Na2CO3 . The reaction takes two steps :

(a) Na2CO3 + HCl NaHCO3 + NaCl Phenolphthalein indicator

(b) NaHCO3 + HCl CO2 + NaCl + H2O Methyl orange indicator

Na2CO3 + 2HCl CO2 + 2 NaCl + H2O

Visual Indicators:

The main purpose of the titration is the determination of the amount of acid or base present in any given solution . This is achieved by locating the equivalence point which gives us the stoichiometric amount of an acid or base chemically equivalent to known amount of a base or an acid .

Indicators are weak organic acids or bases that maintain a color change in definite pH value . The pH value in which color change takes place varies with different indicators therefore we can select an indicator which exhibits the color change at pH close to that obtained at the equivalence point. We mentioned that the change from acid color to alkaline color is not abrupt but small interval of about two pH units involved. Considering undissociated acid indicator as (HIn) , and undissociated base as (InOH) and both having different colors from their ions , the equilibria in aqueous solution will be :

HIn H + + In- and InOH OH - + In -

By applying the law of mass action and considering the activity coefficient as one

(constant),

Kin = and [H +] = =

this equation shows that the actual color change of the indicator is directly related to the hydrogen-ion concentration . Applying the definition of pH , the equation can be written as ,

pH = - log + pKin

For weak base indicator, [OH -] = x Kn , we know that [OH -] =

Substituting for OH -, [H+] = .

Considering acidic form of an indicator as Ina and basic form as Inb, the equilibrium constant is expressed as :

Ina H+ + Inb , KIn = , and [H+] = x KIn

Taking the negative logarithm , pH = pKIn + log . Considering this above equation , the acidic color limit will be :

pH = pKIn - log , in which pH = pKIn – 1

the corresponding alkaline limit will be :

pH = pKIn - log , in which pH = pKIn + 1

therefore the color change interval will be pH = pKIn

Table (2), shows different indicators with their pH range , color change and their preparations.

Indicator

pH - range

Acid color

Base color

Preparation

Methyl orange

3.1 – 4.4

Red

Yellow

0.01% in H2O (Na-form)

Methyl red

4.8 – 6.0

Red

Yellow

0.02g in 60ml EtOH + 40ml H2O

Phenolphthalein

8.0 – 9.6

Colorless

Red

0.05g in 50ml EtOH + 50ml H2O

Bromo cresol purple

5.2 – 6.8

Yellow

Purple

0.1g in 18.5ml of 0.01M NaOH + 225ml H2O

P-nitrophenol

5.6 – 7.6

Colorless

yellow

0.1% in H2O

Bromothymol blue

6.0 – 7.6

Yellow

Blue

0.1g in 16ml of 0.01M NaOH + 225ml H2O

Phenol red

6.4 – 8.0

Yellow

Red

0.1g in 28.2ml of 0.01M NaOH + 225ml H2O

Neutral red

6.8 – 8.0

Red

Orange

0.01g in 50ml EtOH 50ml H2O

Cresol purple

7.6 – 9.2

Yellow

Purple

0.1g in 50ml EtOH + 50ml H2O

Selected Experiments

· Acid – base titrations.

Experiment (1–1), standardization of HCl against sodium carbonate (Na2CO3).

· Titration apparatus

A. Burette (50 ml)

B. Burette Stand

C. Pipette (10 ml)

D. Conical Flask (150 or 250 ml)

E. Methyl orange indicator solution

· Required Solutions.

A. 0.1N potassium hydrogen phthalate (20.3 g/liter)

B. 0.1N anhydrous sodium carbonate (Na2CO3 5.3 g/liter ) .

C. 0.1N HCl ( 8.4ml/liter )

D. 0.1N NaOH ( 4.0 g/liter) .

Procedure: Take 10 ml of the prepared sodium carbonate into 250 ml conical flask and add 2-3 drops of methyl orange ( yellow color will be formed ) . Use unit-volume or graduated pipette in the process .

Fill the burette with the hydrochloric acid ; adjust the volume reading exactly and record it . Remove any traces of air inside the solution . Begin the titration by adding the HCl from the burette onto the carbonate solution carefully and systematically until the end point is close enough . Then make the addition drop-wise until the last drop turns the solution persistent pink color ; and record the values . Repeat the experiment several times until the values are close together within 2-3 units .

Take the average and report as follows:

Initial burette reading

Final burette reading

Vol. reacted, Titer (ml)

0.00

12.20

12.20

12.20

23.50

11.30

23.50

34.20

11.30

Average titer = 11.60

We know that one equivalent of any substance reacts practically or theoretically with one equivalent of the other. Therefore number of equivalents of Na2CO3 reacted = Number of equivalents of HCl reacted. So that, 0.01-liter x 0.1 equiv./liter = 0.0116-liter x NHCl

NHCl = x 0.086N , which is the concentration of the standardized hydrochloric acid .

Experiment (1-2) :

Standardization of sodium hydroxide against potassium hydrogen phthalate(A.R), or against sulfuric acid (optional).

Reagents:

A. Titration apparatus

B. 0.1N NaOH

C. 0.1N Potassium hydrogen phthalate (C8H5KO4) or 0.1N H2SO4

D. Phenolphthalein indicator (0.05 g in 100 ml water/ethanol (1:1) .

Procedure: Transfer 10ml NaOH with a pipette into 250ml conical flask . Rinse the walls of the flask with distilled water . Fill the burette with the sulfuric acid or with the potassium hydrogen phthalate and adjust the volume to specific reading . Begin to add the acid into the base until the end point gets closer ( the color is about to disappear ) . Then begin to add drop-wise until the last drop turns the solution colorless . Tabulate the data and perform the necessary calculations as experiment (1-1) .

Experiment (1 - 3). Determination of calcium carbonate in an impure sample (back titration technique).

Calcium carbonate is an insoluble salt , so definite amount of an excess acid is added to dissolve. After the formation of a clear solution, the acid is back titrated with a base. The sample could be limestone, marble, egg shell etc.

Requirements :

A. Titration apparatus

B. Standard 1M HCl

C. Standardized 0.1N NaOH

D. Methyl orange indicator (0.01 g – soluble M.O in water ).

E. Powdered sample

Weigh accurately 1- 1.5 g . of the carbonate sample and transfer it to 400ml beaker. Add about 20 – 25 ml distilled water, then add 40 ml of 1M HCl; when the effervescence ceases completely , transfer the contents into a 100ml volumetric flask and dilute it to the mark so that you can establish your own dilution factors . Titrate definite aliquots of the diluted solution with sodium hydroxide using methyl orange indicator. calculate the amount of calcium carbonate in the sample

CaCO3 + 2HCl CaCl2 + H2O + CO2

1 mole 2moles

M.wtCaCO3 = 100

Calculations : 100 x 100 x = % by weight of CaCO3 .

Experiment (1-4): Determination of both carbonate and hydroxide in a mixture" commercial caustic soda analysis."

Requirements:

A. Commercial sodium hydroxide

B. Standard 0.1N HCl

C. 0.04M BaCl2 solution

D. Methyl orange and phenolphthalein indicators

In this experiment , total alkali is determined (both , CO32- and OH-) by titrating a portion of the mixture solution with standard acid .The other portion is added an excess of barium chloride solution to precipitate all the carbonate , and without filtering the solution is titrated with standard acid ; this will give the hydroxide volume . Subtracting the second volume from the previous titration volume , will give the volume of the acid that neutralized the carbonate .

BaCl2 + Na2CO3 BaCO3 + 2NaCl

V = volume of the acid using methyl orange ( total alkali )

v = volume of the acid due to phenolphthalein ( hydroxide ) then,

V – 2(V-v) = the volume of the acid that neutralized the hydroxide. and 2(V – v) is the volume of the acid that neutralized the carbonate .

Procedure: Prepare commercial sodium hydroxide solution by accurately weighing about 2.5 g and transfer it quickly into 500ml volumetric flask , dissolve it well with deionized water and fill it up to the mark . Titrate 25ml portion of this solution with standard o.1N HCl using 3drops of methyl orange as an indicator . Repeat the process several times until you get almost the same reading .

Warm another 25ml portion up to 75oC and add 0.04 M BaCl2 solution (quantitatively from burette in a slight excess ). Cool the solution , add 2drops of phenolphthalein, and titrate with 0.1M HCl with constant stirring until the solution turns colorless . Repeat the process several times to get acceptable values .

Keep in mind that 1ml of 1M HCl 0.04 g NaOH , and

1ml of 1M HCl 0.053 g Na2CO3 .

Experiment (1-5): Determination of aspirin in aspirin tablet

Aspirin is an organic compound with carboxylic and an ester group . By hydrolyzing with an alkali , salt of weak acid is formed which can be neutralized with a dilute acid .

C6H4(CO2H)OOCCH3 + NaOH/H2O C6H4(OH)CO2Na + CH3COONa

Aspirin (acetylsalicylic acid) sodium salicylate sodium acetate

The alkali sodium acetate is titrated with dilute sulfuric acid , using phenol red or phenolphthalein as an indicator .

Requirements:

A. Titration apparatus

B. 1M sodium hydroxide

C. 0.05M sulfuric acid

D. Phenol red or phenolphthalein indicator

E. Aspirin tablets

Exact number of aspirin tablets (1.5 – 2.0 g) are accurately weighed and transferred into a conical flask . Add 25ml of 1M NaOH and 25ml of distilled water. Simmer the mixture gently for about 10 – 15 minutes to hydrolyze the tablets. After cooling the mixture , transfer quantitatively to 250ml volumetric flask . Dilute to the mark and mix well by repeated inversions. Titrate aliquots of the reaction mixture with standard 0.05M H2SO4 using phenol red or phenolphthalein .

· Calculate the weight of acetylsalicylic acid .

· Molecular wt. of acetylsalicylic acid = 180.16.

· 1mol of H2SO4 1mol of acetylsalicylic acid.

Experiment (1-6): Determination of available nitrogen in fertilizers.

If ammonium salts are treated with an excess of strong alkali like sodium hydroxide ammonia gas is liberated. The gas is then trapped in excess of dilute acid which can be back titrated with dilute sodium hydroxide.

NH4Cl + NaOH heat NH3 + NaCl + H2O , or NH4+ + OH- NH3 + H2O

Keep in mind that nitrogen compounds are essential nutrients in most fertilizers.

Requirements:

A. Distillation assembly

B. Titration apparatus

C. Standard 0.1N sodium hydroxide

D. Standard 0.1N hydrochloric acid

E. Methyl orange indicator

F. Fertilizer sample (or NH4Cl salt)

Procedure: The distillation apparatus is set up in such a way that the receiver adapter is immersed below the surface of the hydrochloric acid .Weigh accurately definite amount of fertilizer (or NH4Cl salt) , not more than 2.0 g , into 50ml distillation flask . Add 25ml of 2N NaOH solution , few boiling chips and quickly connect to the distillation apparatus. Heat the contents gently to get smooth boiling. Avoid any sucking back from the receiver adapter during boiling . Continue the distillation process until about 10ml of the original solution remains ( 30 – 40 min) .

When the distillation is over , disconnect the apparatus from top before removing the Bunsen burner or any other heating system . Wash any residues in the condenser and receiver adapter into a 150ml conical flask and titrate against .0.1N NaOH using 3drops of methyl orange indicator.

Calculate:

A. The weight of ammonia distilled.

B. Percent of available nitrogen in the fertilizer.

· Precipitation Titrations

Precipitation titrations are among the oldest analytical methods employed in chemical analysis. Up to now they are applied in certain volumetric analysis such as the determination of silver , chlorides , bromides , iodides and thiocyanates. In this method , certain organic indicators are used ; these indicators have the property of adsorption or desorption on the solid materials formed during the precipitation . This adsorption and desorption process takes places near the equivalence point accompanied by change in the color ; the procedure is called Fagan’s method as an honour for Polish scientist Kazimierz Fagan’s ; among these indicators are fluorescein and its derivatives . Since most of these procedures involve the application of silver nitrate solutions they are called argentometric methods.

We have mentioned that the most important precipitation titration utilizes silver nitrate as the reagent for the process ; therefore in these titrations , we will theoretically confine our discussions to argentometric reactions .Before we launch the practical aspect of this titration , let us consider certain important concepts relating the solubility of the salt AgCl which forms during the titration process. Silver chloride is sparingly soluble in water, and considering its (AgCl) saturated solution the solubility equilibrium will be represented as:

AgCl(s) Ag (aq) + Cl -(aq)

Ksp =

This is an example of heterogeneous equilibrium so the concentration of the solid part (AgCl) remains constant in the solution and can be taken as unity (1) so that :

Ksp = [Ag+][Cl-]

This molar product of Ag+ and Cl - at equilibrium is called solubility product . In expressing the solubility product of the ions involved , the stoichiometric coefficients are raised .

Ag2S 2Ag+ + S2- , Ksp = [Ag+]2[S2-]

نتيجة بحث الصور عن ‪table solubility product constant‬‏

Table 2.2 , shows the solubility product (Ksp) of certain slightly soluble salts at 25oC .

It must be understood that the smaller the value , the less soluble salt in water . We can determine the Ksp from the molar concentration at equilibrium and vice versa . Ksp is comparatively easy to calculate with respect to other equilibrium calculations .

Example 1.

An aqueous saturated solution of silver carbonate was analysed and found to contain 8 mg of the salt dissolved in 250 ml. Calculate the Ksp of the Ag2CO3 .

Solution:

Ag2CO3 2 Ag+ + CO32- ,

it shows that in every one mole of silver chloride dissolved two moles of silver ions and one mole of carbonate ion will be produced.

Ksp = [2Ag+]2[CO32-] , the molecular weight of Ag2CO3 = 275 . 8

Amount of mg of Ag2CO3 = = 32 mg /l

Molar solution of Ag2CO3 = = 1.16 × 10 -4 M

Ksp = [2Ag+]2[CO32-] = [2 X 1.16 x 10 – 4]2 [1.16 X 10 – 4] = 6.4X 10 – 12

We said that the reverse is true , that is we can calculate the molar concentration from the solubility product , Ksp .

Example 2. Calculate the molar solubility of silver sulphate solution whose solubility product at 25 0C is 1.4 X 10 -5 .

Ag2SO4 2 Ag + + SO42-

Ksp = [2Ag +] 2[SO42-] = 1.4 X 10 -5

Let us assume that x moles per litter dissolved in the solution , so [Ag +] = 2 x, [SO42] = x So substituting these values into the equation , Ksp = [2Ag +] 2[SO42-] = 1.4 x 10 -5

(2x)2(x) = 1.4 X 10 -5 and 4 X3 = 1.4 X 10 -5 , X = = 1.5 X 10-2

[SO4-2] = [X] = 1.5 X 10-2 , [Ag+] = [2X] = 3.0 X 10-2

Determination of end-point in precipitation titration

There are different methods to determine the end-points of the precipitation titrations. One of these methods is the application of chemical indicator which can be presented in the following ways:

a. Appearance of colored precipitate; good example of this end – point detection is the application of Mohr's method of titration . In this method , the indicator used is dilute solution of potassium chromate (K2CrO4) . Silver chloride is less soluble than potassium chromate and for the silver chloride to precipitate first in the process , its initial concentration must be high .

[Ag+][Cl -] = 1.6 X 10-10

[2Ag+]2[CrO42-] = 1.7 X 10-12

Based on their solubility product , more dilute solution of K2CrO4 must be used to give red sharp end-point . It must be noted that concentrations of chromate ion as large as 0.006 M masks the red color of the Ag2CrO4 due to persistence of the color of K2CrO4 .

The pH of the solution in which this reaction takes place must be considered. In more acidic solution , dichromate is formed which is more soluble than chromate and may need large silver concentration .

2CrO42- + 2H+ Cr2O72-+ H2O

In strong alkaline solution , the silver may precipitate as silver hydroxide which changes into silver oxide .

2Ag+ + 2OH - 2AgOH(s) Ag2O(s) + H2O

b. Formation of soluble colored complex .

In Volhard method, standard solution of thiocyanate may be used for silver ion titration .

Ag+ + SCN - AgSCN

The indicator for this titration is Fe3+ solution , so that the last excess drop turns the solution red due to complex formation ,

Fe 3+ + SCN - Fe(SCN)2+

To avoid the precipitation of Fe3+ hydroxide or hydrated oxide , the titration must be carried out in acidic solution .

c. Adsorption indicators.

Fajan's method of chloride titration utilizes specific indicators called adsorption indicators . An adsorption indicator is an organic substance that has the property of being adsorbed on or desorbed from the surface of the precipitate . The adsorption and desorption takes places theoretically near the equivalence point accompanied with color change. Titrations that employ adsorption indicators are generally accurate , fast and reliable

Fluorescein is typical organic dye used for the titration of chloride ions with silver nitrate .

نتيجة بحث الصور عن ‪fluorescein‬‏ نتيجة بحث الصور عن ‪fluorescein‬‏

yellow red Na-form

derivatives of fluorescein like 2,7- Dichlorofluorescein is preferably used in Fajan's method of chloride determination .

نتيجة بحث الصور عن ‪dichlorofluorescein‬‏

2.7- Dichlorofluorescein

Experiment (2-1) Chloride determination by Fajans method.

Reagents:

A. 0.1N AgNO3 (standardized) 17.0 g dissolved in H2O , dilute to liter .

B. Dichlorofluorescein solution ( 0.1% ethanol solution ) .

C. Dextrin (D-glucose polymer) .

Procedure: Prepare a suitable chloride solution from an analytical reagent salt in one liter volumetric flask . Take 50 ml aliquot of the solution into 250 ml conical flask ; add five drops of the indicator and 0.1 g of dextrin .

Titrate against the standardized AgNO3 solution from the burette until permanent pink color is obtained . Repeat the titration three times and report the chloride content in mg/l .

Experiment (2-2): Determination of chloride by Mohr method

The most significant idea in this method is the application of K2CrO4 solution as an indicator to detect the end point . Even though both AgCl and Ag2CrO4 are insoluble salts , the solubility product of Ag2CrO4 is far more greater than that of AgCl ; thus AgCl precipitates first during the titration process and first silver ions after the total precipitation of of chloride ion react with the chromate ions forming brick-red color as an end point .

Reagents :

A. Standard AgNO3 solution (0.1 N)

B. 0.25 M K2CrO4 solution

Procedure: Take 50 ml aliquot of the prepared solution in the preceding experiment (unknown) into 250 ml conical flask ; add 2 ml of the K2CrO4 solution. Start the titration by adding AgNO3 solution from the burette to the chloride solution until the permanent brick-red color appears . repeat the process three times and report as before .

Experiment (2-3): Determination of Ag in a solution (Volhard’s method)

In this method , the amount of silver in a sample is determined by titrating against standardized KSCN solution . Ferric alum (ammonium iron(III) sulfate ) is used to indicate the end point . In this titration :

Ag+ + SCN - AgSCN

And Fe3+ + SCN - Fe(SCN)2+ (reddish-brown)

Reagents:

a. 0.1 N KSCN solution

b. Ferric alum , saturation solution containing 6M HNO3 free of nitrogen oxides.

c. Solution containing Ag+ (unknown) .

Transfer 50 ml aliquot of the sample solution into 250 ml conical flask . Add 2 ml of the ferric alum solution . Start the titration by adding KSCN solution from the burette with constant vigorous agitation until reddish brown color is formed . Report the result as before and calculate the of Ag .

3. Complexometric titrations – EDTA titrations

The most important complexing agent in terms of stoichiometric titrations is the EDTA (Ethylenediaminetetraacetic acid ) . This organic complexing agent has more than two chelating sites . Titrations involving these chelating agents are called Complexometric titrations. EDTA forms stable complexing compound with many metal cations at different pH values . Calcium and magnesium ions are some of the metal cations that are relatively stable EDTA complex ; their formation constants are too close together so that they can be titrated together using Eriochrome black T (EBT) as an end point indicator , so this titration can be used to determine the hardness of water .

EBT forms too weak complex with calcium giving no sharp end point , therefore small known amount of Mg2+ is added to the Ca2+ solution which can be corrected for it at the end of the titration by using blank titration with the same amount of Mg2+.

We have already explained in the theoretical section that different metal ions react with electron pair donor and form either a coordination compound or complex ions . This reaction generally involves the replacement of the solvent molecules coordinated with the metal ion with other donors . These replacing donors are called ligands and must have at least one pair of unshared electrons ; they can be a charged ion or neutral molecule ; they can also be organic or inorganic species and the complex formed can be ion or neutral like , Cl -, NH3 , H2O , NH2-CH2-COO- etc.

In complex formation , the reaction occurs in steps and step-wise equilibria is formed . Let us take the formation of the complex ML n .

M + L1 ML1 K1 =

M + 2L ML2 K2 =

M + nL MLn Kn =

Each step has its own step-wise stability constant ; the overall stability constant K is related to the wise constants .

K = K1 . k2 . Kn

نتيجة بحث الصور عن ‪edta‬‏This relationship is valid only if there no insoluble product formed. One of the most widely used aminopolycarboxylic acid is the Ethylenediaminetetraacetic acid , EDTA. For simplification , we can take H4Y instead of EDTA and Na2H2Y instead of the disodium form . The substance is weak organic acid which undergoes four stability constants .

نتيجة بحث الصور عن ‪edta‬‏

H4Y H3Y - + H+ K1 = 1.02 X 10 -2

H3Y 3- H2Y 2- + H + K2 = 2.14 X 10 -3

H2Y 2- HY 3- + H + K3 = 6.92 X 10 -7

HY 3- Y 4- + H + K4 = 5.5 X 10 -11

The K values show the ease of the hydrogen ion to be released .The first two values are much greater than the others. The sodium salt is generally used in titrations for determining different metal cations . The important feature of the EDTA titrations is that it combines with all metal ions in ratio of 1:1 and forms stable chelates suitable for reproducible volumetric analysis .

· Types of EDTA titrations:

1. Direct titrations: In this method direction titration is made between the metal ion and standard solution of EDTA . The metal ion solution is usually buffered with NH4+/NH4Cl at pH – 10 .

In certain cases it is very important to prevent any metal hydroxide precipitation by using additional complexing agents like, citrate , tartrate or other organic chelating agent like , 2,2',2"- Nitrilotriethanol (Triethanolamine). The equivalence point is accompanied with change in color of the metal indicator . The end-point can be determined instrumentally by conductometric , amperometric or spectrophotometric methods .

2. Back titration methods: Sometimes it is not possible or difficult to determine the metal by direct titration with EDTA . There are several reasons for this problem ; sometimes the pH necessary for the titration may precipitate the metal , other times undesirable complex may be formed or no suitable indicator is available . In these situations back titration method is applied in which excess of EDTA solution is added to the metal solution and the pH of the resulting solution is adjusted to the desired value , then the excess of the EDTA is titrated with standardized metal ion solution.

نتيجة بحث الصور عن ‪eriochrome black t‬‏To give an example , the reaction between Al3+ and EDTA is very slow , so direct titration is very difficult . The problem is overcome by adding excess amount of EDTA to the Al3+ solution and let the mixture stand still for about 35 minutes . Then the excess EDTA is titrated with standard Zn2+ solution . The Al3+ is calculated from the amount of EDTA consumed.

3. Displacement titration: In this system, metal ions that do not form stable complex (or do not react) with the metal indicator has to be substituted for metal ions which form more stable EDTA complex . Usually this happens in the case of magnesium or zinc complex .

MgY2- + M MY2- + Mg2+

The displaced magnesium is back titrated with standard EDTA.

Metal ion indicators (EDTA Titrations)

The result of EDTA titration depends on precisely locating the end-point ; this is usually achieved by utilizing specific indicators called metal ion indicators. These indicators are organic dyes that chelate with the metal ion in a given range of concentrations, forming specific color . The intensity of the color is discernible within that concentration range (10-6 to 10-7) . The most widely used metal indicator is the Eriochrome Black T

Eriochrome Black T

The sulfonic acid functional group ionizes completely in aqueous solution releasing hydrogen ions but the phenolic group partially ionizes .

H2In- + H2O HIn2- + H3+O K1 = 5 x 10 -7

red blue

HIn2- + H2O In3- + H3+O K2 = 2.8 x 10-12

Most of the EBT metal complexes have generally red color and to observe the color change during the process , the pH must be adjusted to 7 and above so that free indicator color (blue} becomes clear. At the end-point:

MIn + HY3- HIn2- + MY2-

The advantage of the EBT is that it complexes with more than twenty different ions, but only certain number of these metal ions form complexes stable enough to detect the end-point.

Experiment (3-1): Determination of water hardness using EDTA.

Reagents:

A. 0.01 M EDTA (3.8 g of Na2EDTA , dried , purified in 1liter).

B. 6 M HCl

C. 6 M NaOH

D. Methyl red indicator

E. Eriochrome black T (100 mg dissolved in a mixture of 15 ml triethanolamine and 5 ml absolute ethanol).

F. Buffer solution pH-10 (570 ml conc. NH3 + 70 g NH4Cl all diluted to 1liter).

Procedure:

Take accurately measured aliquot of water (not less than 100 ml) , add 0.5 ml of 6 M HCl and boil for few minutes to get rid of CO2 . After cooling the sample add few drops of methyl red and neutralize with NaOH . Add 2 ml of the buffer solution followed by 5 drops of EBT or calmagite indicator . Start titrating the sample with the standard EDTA from the burette until the red color changes to blue . Repeat the process three times and report the result as mgs of CaCO3/liter.

Experiment (3-2a) Determination of magnesium by titrating with EDTA .

Reagents:

1. Standard EDTA (0.01M)

2. Buffer pH-10 (570 ml conc. NH3 + 70 g NH4Cl diluted to liter)

3. Eriochrome Black T indicator

Take prepared sample solution of Mg2+ into 250 ml volumetric flask and dilute to the mark . Take 50 ml of the diluted sample solution into 250 ml conical flask add 2 ml of the buffer solution and 5 drops of the indicator . Titrate with 0.01M Na2H2Y until the color changes from red to blue . Repeat the titration two more times . Estimate the Mg in milligrams/ liter or ppm .

Experiment (3-2b): Determination of Ca by Complexometric back titration

Reagents:

1. EDTA (0.01M , standardized)

2. 0.01M MgSO4 (1.3g MgSO4 . 7H2O in 500 ml distilled water)

3. 6M NaOH

4. 6M HCl

5. Methyl red indicator solution

6. Eriochrome Black T indicator solution (EBT)

Take 25 ml of 0.01M Ca2+ solution into 250 ml conical flasks , add 25 ml distilled water , 2ml buffer solution pH-10 , 5 drops of EBT indicator calamite . Run excess of Na2H2Y from the burette and record the volume . Titrate the excess EDTA with 0.01M MgSO4 solution until the color changes from blue to red . Calculate the amount of Ca2+ in mg/liter .

1 mole of EDTA 1 mole Ca2+

4. Redox – Titrations

We have already seen that oxidation – reduction systems deal with reactions in which loss and gain of electrons are involved . In these reactions , titrimetric procedures can be employed to determine quantitatively one of the reducing or oxidizing agents in a given sample . Like acid – base titrations , we employ indicating system to denote the equivalence point during the titration . The indicating technique as we mentioned before, can be a change in color , redox potential or self – indicating system .

In this section we will briefly discuss about common oxidizing agents in oxidation-reduction titrations such as potassium permanganate , potassium dichromate , cerium (IV) sulfate , and iodine despite that there are many others that fall into this category .

A. Potassium permanganate

This substance is strong oxidizing agent and takes very important role in oxidation-reduction titrations . It is applied as an oxidizing agent in many organic and inorganic reactions in both acidic and basic solutions. In acidic solution , sulfuric acid is usually used since it has no effect upon permanganate in dilute solutions unlike the hydrochloric acid which may reduce the chloride ion and decrease the oxidizing power.

2MnO4- + 16H+ + 10Cl - 2Mn2+ + 5Cl2 + 8H2O

The reduction half-reaction of the permanganate in acidic solution can be written as :

MnO4- + 8H+ + 5e Mn2+ + 4H2O

Since the standard potential is 1.51 volts in acidic solution , it is considered as powerful oxidizing agent. Another advantage of the MnO4- is that there is no need for an indicator in colorless solutions . More important feature of the permanganate is that it can be applied in strong alkaline solutions .

MnO4- + 2H2O + 3e MnO2 + 4OH –

Potassium permanganate cannot be taken as primary standard because it is rarely obtained in pure form since there are always traces of manganese dioxide and remains of organic substances from the distilled water . These substances may reduce the permanganate to manganese dioxide which catalyzes the self-decomposition of the permanganate solution on standing for long time .

MnO4- + 2H2O MnO2 + 3O2 + 4OH –

At the same time any traces of Mn2+ affects the stability of the solution and careful steps has to be followed when preparing the MnO4- solution ; the solution must be covered from unnecessary exposure to bright light. Standardization of the potassium permanganate solution is accomplished by titrating with a standard solution like sodium oxalate in acidic solution .

2MnO4 + 5H2C2O4 + 6H+ 2Mn2+ + 10CO2 + 8H2O

The oxalate solution must be heated to about 90 oC .

· Preparation of 0.1N KMnO4.

نتيجة بحث الصور عن ‪sodium diphenylamine sulfonate‬‏Weigh out accurately 3.16g of A.R potassium permanganate into one liter volumetric flask and dissolve it in distilled water by shaking well during the dissolution process. Fill it to the mark and store it in a dark or in diffuse light when is not in use .

B. Potassium dichromate (k2Cr2O7) .

Potassium dichromate is not as powerful oxidizing agent as potassium permanganate , yet it has several advantages over the latter . Potassium dichromate can be obtained in a pure form . It is least affected by direct light , heat and water vapor in the atmosphere . Another advantage is that it is used in acidic solution and quickly reduces to green Cr3+ .

Cr2O72- + 14H+ + 6e 2Cr3+ + 7H2O

as a result it is considered as excellent primary standard . Dichromate solution easily oxidizes the organic compounds in comparison to permanganate solution . The appearance of the green color due to formation of Cr3+ makes the end-point almost impossible to be visually ascertained ; so redox indicator is employed to give clear and distinct color change . Suitable indicators for this titration are A. N-phenyl anthranilic acid (C13H11O2 , 0.1% in 5x10-3N NaOH). نتيجة بحث الصور عن ‪n phenylanthranilic acid indicator preparation‬‏N-phenyl anthranilic acid

B. another important indicator is sodium diphenylamine sulfonate (C12H10NNaO3S , 0.2% , w/v , H2O). sodium diphenylamine sulfonate

Preparation of 0.1N K2Cr2O7.

Weight accurately about 4.9 g. of finely powdered and well – dried K2Cr2O7 into 1-liter volumetric flask using suitable funnel to avoid any loss of the salt . Dissolve the salt in distilled water and dilute to the mark . This solution can be standardized with other reducing agent like iron(II) salt .

Preparation of 0.1N Iron(II) sulfate.

This solution can be prepared by accurately weighing 27. 8 g. of ( A.R) FeSO4.7H2O or 39.21g. of (A.R) Iron(II) ammonium sulfate (FeSO4.(NH4)2SO4.6H2O) into 1-liter volumetric flask . Dissolve the salt in distilled water and dilute to the mark.

C. Cerium (IV) Sulfate .

Solution of cerium (IV) sulfate in sulfuric acid is powerful oxidizing agent . It is more powerful than the potassium permanganate and can be a substitute for it in most cases . Cerium (IV) can only be used in acidic solution of 5N and more. There are considerable advantages of Ce(IV) sulfate over other oxidizing agents as a reliable primary standard .

Ce(IV) sulfate solution is indefinitely stable and not affected by exposure to direct light or heat unlike the permanganate solution .

Ce(IV) sulfate solution does not oxidize chloride ions , even in higher concentrations contrary to the situation in potassium permanganate .

There are no variable oxidation states when reacting with reducing agents so that the equivalent weight and molar weight are the same .

Ce4+ + 1e Ce3+

These properties besides others makes the Ce(IV) solution preferable substitute for permanganate solution . different Ce(IV) salts are commercially available in high standard of purity making it suitable for preparing excellent primary standard solution .

Preparation of 0.1N Ce(IV) sulfate.

66 g of ammonium ceric sulfate [Ce(SO4)2 . 2(NH4)2SO4. 2H2O] is accurately weighed and transferred into a mixture of 30 ml of conc. H2SO4 and 30 ml of distilled water . Heat gently to dissolve , cool the solution and filter it through fine – porosity sintered glass crucible ; dilute to liter and mix thoroughly . The solution can be standardized with 0.1N sodium oxalate or with ammonium Fe (II) sulfate using Ferroin ( 1,10-phenanthroline Fe (II) sulfate ) as indicator .

D. Iodimetry titration.

Titration analysis relating iodine solutions is called Iodometric titrations . Saturated aqueous solution of iodine is about 1 x 10-2M at room temperature , so to achieve higher concentrations of iodine solution we have to introduce external iodide ions to form soluble tri-iodide complex .

I2(s) + I - I3- K = 7.1 x 10 2

This external iodide ions comes from potassium iodide salt , so in iodimetry titrations we have to employ concentrated iodide solution .

I2 + 2S2O32- 2I - + S4O62- Or

I3- + 2S2O32- 3I - + S4O62-

Iodine is weak oxidizing agent and employed in the titrations of certain strong reducing agents . Iodine solution is not stable solution and needs re-standardization from time to time ; one of the reasons for instability is that iodine crystals are volatile (sublime) so there may be certain loss during the preparation of the solution. Also , iodine slowly attacks the stoppers other than glass . Another reason may be a change in the solution concentration due to possible air oxidation .

4I - + O2 2I2 + 2H2O

The reaction is catalyzed by light , heat and acidic traces . The titration reaction is sufficient to be self-indicating . Aqueous solution of iodine has yellowish brown color and disappearance of this color during the titration can serve as the end-point of the titration. Although the system can serve as self – indicator , starch solution is used to give sharp end – point . Starch gives an intense blue color with I2 even in very low concentrations .

Preparation of starch:

Take about 2g of soluble starch into about one liter of water . Boil the mixture until clear solution is obtained . Cool and store in stoppered bottle .

Preparation of 0.1N iodine solution

Weigh 20 g of potassium iodide (free of iodate) , into 40ml of distilled water in a glass stoppered volumetric flask . Weigh out about 12.7g of iodine (use rough balance to avoid iodine vapor) and transfer it to potassium iodide solution . Stopper the glass and shake well until the iodine is dissolved and fill to the mark .

Experiment (4-1) Determination of Iron(II) by titrating with K2Cr2O7 solution .

Requirements :

A. Titration apparatus

B. 0.1N K2Cr2O7 (A.R)

C. 0.1N Iron(II) sulfate or ammonium Iron(II) sulfate .

D. Sodium diphenylamine sulfonate (1% solution in 0.005N NaOH ).

E. Ortho-phosphoric acid (concentrated).

Titrate 25.0ml of the acidified Iron(II) sulfate solution with standard 0.1N potassium dichromate solution using sodium diphenylamine sulfonate solution (0.5ml) and 2.5ml concentrated phosphoric acid . Continue adding the dichromate solution slowly with constant stirring until the green color changes into a grey-green . Then begin to add the dichromate drop wise until the first blue – violet appears permanently . Repeat the process three times , report the result as usual and calculate the amount of Iron .