Redox Titration. Redox Reactions Redox stands for Reduction - Oxidation

Redox Titration Dr. Manisha Jain Asst. professor (Inorganic chemistry) Department Of Chemistry, Acharya Narendra Dev College (University of Delhi). De...
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Redox Titration Dr. Manisha Jain Asst. professor (Inorganic chemistry) Department Of Chemistry, Acharya Narendra Dev College (University of Delhi). Delhi, India. Redox Reactions Redox stands for Reduction - Oxidation These titrations involve the titration of an oxidizing agent (or oxidant) with a reducing agent (or reductant) or vice versa. There must be a sufficiently large difference between the oxidizing and reducing capabilities of theses agents for the reaction to undergo completion with a sharp end point. Oxidation process involves loss of electrons while reduction process involves gain of electrons. Thus an oxidizing agent is one which accepts electrons while a reducing agent is one which loses the electrons. An oxidizing agent oxidizes the other substance by stripping off electrons from it. A reducing agent reduces the other substance by donating electrons to it. Oxidation and reduction reactions always occur simultaneously. One can not take place in isolation from the other. During a redox reaction the oxidizing agent itself undergoes reduction while the reducing agent undergoes its oxidation. Thus reduction reaction is represented as Oxidant + ne = Reductant Oxidation reaction is represented as Reductant = Oxidant + ne Oxidation of a substance leads to increase in its oxidation number while Reduction of a substance lead to decrease in the oxidation number. The relative tendency to accept or lose electrons by any reagent is measured in terms of their standard reduction potential values. Relative values are only obtained by comparison with the Normal Hydrogen electrode potential (SEP) when the hydrogen gas at 1 atm. pressure is in contact with a piece of platinum and with hydrogen ion concentration at 1M. The standard electrode potential enable us to predict which ions will oxidize or reduce other ions. A reversible redox system may be represented as Oxidant + ne = Reductant Or Oxd + ne = Red

The electrode potential which is established when an inert or unattackable electrode is immersed in a solution containing both the oxidant and the reductant is given by the expression ET = E0 + RT/nF ln [oxd]/[Red] Where ET is the observed potential of the redox electrode at temperature T , E 0 is the standard reduction potential, n the number of electrons gained by the oxidant in being converted to the reductant. The Electrochemical series enlists a number of systems according to decreasing standard reduction potentials at 25 0C. The most powerful oxidizing agents lie at the top of electrochemical series ( High positive E0red values) and the most powerful reducing agents are present at the bottom (High negative E0red values). Some of the commonly used oxidizing and reducing agents in the redox titrations are – Oxidizing agents i. KMnO4 in presence of dil H2SO4 MnO4

¯

+ 8H+ + 5e-



Mn2+

+ 4H2O

E°red = +1.52V

ii. K2Cr2O7 in dil. H2SO4 a moderately strong oxidizing agent; oxidizing ability depends strongly on pH, decreasing rapidly as solution becomes more neutral Cr2O72– + 14H+ + 6e– → 2Cr3+ + 7H2O iii.

E°red = +1.33V

Iodine solution I2 + 2 e- = 2I–

E°red = +0.54V

Reducing agents i.

Mohr’s salt

FeSO4.(NH4) 2SO4.6H2O Fe2+

ii.

Oxalic acid



Fe3+

+ e-

H2C2O4.2H2O C2O42- = 2CO2 + 2 e-

iii.

E°red = +0.77V

E°red = +0.77V

Sodium thiosulphate Na2S2O3.5H2O 2S2O32- = S4O62- + 2 e-

E°red = +0.08V

There is no universal oxidizing agent which can be titrated against every reducing agent and vice-versa. Hence, the choice of an oxidizing agent to be used against a particular reducing agent depends upon the reaction conditions and standard reduction potential of the oxidizing agent. pH dependence of oxidizing behaviour It is important to note that for many oxidants the pH of the medium is of great importance and hence their oxidizing strength may vary depending on the medium in which its reaction is studied. For example potassium permanganate is oxidizing agent in all three mediums, scid, alkaline and neutral. However it is strongest in acidic medium. a. Strongly alkaline medium MnO4 + e- → MnO42-

E°red = +0.56V

¯

Permanganate ion

Manganate ion

Equivalent weight of KMnO4 = b. Neutral medium MnO 4 + 2H2O ¯

+

Molecular weight ---------------------1

3e --



MnO 2↓ + 4OH-

E°red = +1.23V

Manganese dioxide ppt.

Equivalent weight of KMnO4= c. Acidic medium MnO4 + 8H+ + 5e¯

Molecular weight ---------------------3 Mn2+



+

4H2O

E°red = +1.51V

Manganous ion

Equivalent weight of KMnO4 =

Molecular weight -------------------------5

Redox Indicators A redox indicator should be such that it produces a sudden change in the electrode potential in the vicinity of the equivalence point during a redox titration. This is possible when the indicator itself is redox active i.e., capable of undergoing oxidation or reduction process which is a reversible one. The oxidized and reduced form of the indicator should have a contrast difference in the colours. Inoxd + ne = Inred At potential E, the ratio of the concentration of two forms is given by the Nernst equation ET = E0 + RT/nF ln [Inoxd]/[IRed] Type of Redox Indicators

Self Indicators Many a times the titrant itself may be so strongly coloured that after the equivalence point, a single drop of the titrant produces an intense colour in the reaction mixture. e.g. potassium permanganate. Such Indicators are called self indicators. Self indicators generally are strongly coloured as a result of charge transfer transitions in them. Internal indicator Such indicators are added into the reaction mixtures Such indicators always have reduction potential values lower than the analyte system so that they react with the titrant only when whole of the analyte has been consumed, producing a readily detectable color change. External indicator In case a suitable redox indicator is not available for a given system, an indicator may be employed which will indicate the completion of reaction by physically or chemically reacting with the analyte (not through redox reaction). This reaction between indicator and the analyte may sometimes be an irreversible one and in some cases may even lead to precipitation. In those case indicators are not added to the reaction mixture on the whole, rather used externally on a grooved tile. Such indicators are called external indicators. Potentiometric Methods The most accurate method to judge the completion of redox titration is however potentiometric method which deals with the measurement of e.m.f. between a reference electrode and the indicator (redox) electrode during the stages of redox titration. This method infact tells us the equivalence point of the reaction and not the end point Titration of Potassium permanganate Vs Oxalic acid Potassium permanganate is not a primary standard. It is not easily obtained in perfectly pure form which is completely free from manganese dioxide. Also the aqueous solutions are not stable for long as the ordinary distilled water usually contains reducing substances (traces of organic matter, etc.) which react with permanganate resulting in precipitation of MnO2. This reaction is catalysed by light. Presence of manganese dioxide is totally undesirable as it catalyses the auto decomposition of the permanganate ion on standing: 4MnO

¯

4

+ 2H2O → 4MnO2 + 3O2 + 4 OH

¯

Permanganate solution decompose in presence of manganous ions also: 2MnO

¯

4

+ 3 Mn2+ + 2H2O = 5MnO2 + 4H+

The reaction is quite slow in acidic medium but attains a fast rate in neutral solutions. Thus Permanganate solution can not be considered as a primary standard solution by simply making the calculations based on the amount of permanganate weighed and

dissolved. The permanganate solutions once prepared are either left for a day or two or the freshly prepared solutions are boiled for an hour. In both the cases the resulting solution is filtered through sintered glass crucibles. The clear solutions so obtained are then standardized by titrating with a primary standard solution Permanganate ion in the acidic medium is a very strong oxidizing agent MnO4 + 8H+ + 5e¯



permanganate

Mn2+

+ 4H2O

E°red = +1.51V

Manganous ion

The acidity is introduced by addition of dil. H 2SO4 only and no other acid. Conc. H2SO4 , Nitric acid (both conc. and dilute) can’t be used as these are oxidizing agent and may interfere in the reaction of permanganate depending on the reductant’s strength. In the hydrochloric acid permanganate can oxidize chloride to chlorine, which can be a source of positive errors as permanganate is consumed in this reaction. (E°red Cl2/Cl-)= +1.36V) 2MnO

¯

4

+ 16 H+ + 10Cl (aq) →

2Mn2+ + 5Cl2 + 8 H2O (l)

-

Acetic acid is too weak an acid to provide the desired acidity in the solution. Permanganate acts as self indicator. Since MnO4– is intense purple while Mn2+ is colourless, the reaction mixture at equivalence point is colourless and even a single drop of the permanganate would impart sufficient pink colour to the solution acting as self indicator. This also leads to precaution that while titrating a reducing agent with KMnO4 , constant stirring of reaction mixture is must otherwise local reduction in acidity may lead to precipitation of MnO2 according to this reaction. The reducing agent in the titration to be discussed is oxalic acid here. The composition of it is H2C2O4.2H2O. Inspite of being a dehydrate it is a good primary standard as its composition is unchanged during storage or weighing. The reaction between oxalic acid and potassium permanganate can be represented as: 2KMnO4 + 5 H2C2O4 + 3H2SO4

= 2MnSO4 + K2SO4 + 10 CO2 + 8H2O

In ionic form the reaction can be represented as: 2MnO4¯ + 10 C2O42¯+ 16H+ = 2Mn2+ + 10 CO2 + 8H2O This redox reaction can be split apart in two parts- one showing the oxidation and the other reduction 2MnO4¯ + 10e¯ + 16H+ = 2Mn2+ + 8H2O + VII

+II

reduction (Oxidation number has decreased)

5C2O42¯ = 10 CO2 + 10e¯ +II

oxidation

+IV

(Oxidation number has increased) This titration is carried out in warm conditions (temperature about 60 oC). The reaction at room temperature is slow because of the equilibrium nature of this reaction. CO 2 is highly soluble in water and thus heating removes all dissolved carbondioxide out of the solution driving the reaction in forward direction. Also at low temperature, the reduction of permanganate may not be complete producing Mn(III) (in the form [Mn(C 2O4)3]3-). The formation of this species introduce errors in titrations as no. of electrons utilized here are different as compared to production of Mn2+. This complex however breaks on heating and hence reaction proceeds smoothly. While noting the burette readings, it should be taken into account that the solution is so intensely coloured that the lower meniscus of the solution may not be clear. Thus for permanganate titrations the upper meniscus in the burette is noted. Another important redox titration is titration of Potassium dichromate with Mohr’s salt. Potassium dichromate is also a very strong oxidizing agent (E°red = +1.33V) . However it is not as strong oxidizing agent as permanganate is (E°red = +1.51V).Still it is widely used in redox titrations because of several advantages over permanganate. Unlike potassium permanganate, potassium dichromate is available in high purity and is highly stable upto its melting point. Its aqueous solutions are not attacked by oxidisable impurities like rubber or any other organic matter and thus composition of aqueous solution does not change on keeping. The aqueous solutions are quite stable towards light. It is thus an excellent primary standard and its standard solutions can be prepared by direct weighing of an amount of it and dissolving in a known volume of distilled water. Potassium dichromate acts as oxidizing agent in acidic medium only: The neutral aqueous solution of Potassium dichromate is 1:1 equilibrium mixture of dichromate and chromate, a consequence of hydrolysis of dichromate ions. Cr2O72– + H2O = Orange

2 CrO42– + 2H+ yellow

Chromate ions are weaker oxidizing agent than dichromate. Thus oxidizing strength of dichromate is reduced in neutral solution. The above hydrolysis reaction however can be reversed by adding acid to the solution and this explains the necessity of acidic medium for the reaction. Also the reduction reaction of dichromate can be represented as: Cr2O72– + 14H+ + 6e– → 2Cr3+ + 7H2O

E°red = +1.33V

+VI Orange

+III green

This reaction clearly shows the involvement of H+ in the reduction half reaction. The medium is generally acidified with dil. H2SO4. Unlike permanganate case, cold HCl can be used here for acidifying the reaction mixture provided the acid conc. Does not increase beyond 1 – 2 M. Though the dichromate solutions are intensely orange coloured solutions and a single drop of it imparts yellow colour to a colourless solution, it can’t be used as a self indicator like KMnO4. This is because its reduction product (Cr3+) is green which hinders in the visual detection of end point by observing dichromate colour. Thus an indicator is must in this titration. The indicator should be redox active and must be properly chosen keeping in mind the electrode potential values of the reducing agent being titrated with dichromate. Suitable indicators for dichromate titrations are Diphenylamine (specifically sodium diphenylamine sulphonate)(0.2%aq. Soln) in presence of orthophosphoric acid, and N-phenylanthranilic acid (0.1% soln. in 0.005 M NaOH). Mohr’s Salt The reducing agent used in this titration is Mohr’s salt which is a double salt. Its composition is FeSO4.(NH4) 2SO4.6H2O. The redox active species in this compound is Fe2+ whose oxidation can be represented as: Fe2+



Fe3+

+ e-

E°red = +0.77V

Mohr’s salt is a primary standard and therefore its standard solution can be easily prepared by direct weighing a known quantity of it and dissolving in a known volume of distilled water. However The Ferrous ions in aqueous solutions are highly susceptible towards hydrolysis resulting in precipitation of ferrous hydroxide. [Fe(H2O)6]2+ = [Fe(H2O)5(OH) ]2+ + H+ Aqua ion ↓↑ Fe(OH)2 + 4 H2O + H+ Brown ppt.

The reaction sequence however can be reversed by adding acid. Hence while preparing standard solution of Mohr’s salt it is necessary to add about 1 ml of dil. H2SO4 to the measuring flask. Simple Ferrous sulphate is not a primary standard and hence can never be used for preparing the standard solution. The solid sample of ferrous sulphate contains a high percentage of ferric ions produced by aerial oxidation of the ferrous ions. Thus composition is not accurately known. The reaction between Potassium dichromate and Mohr’s salt can be represented as:

K2Cr2O7 + 6FeSO4 + 7H2SO4 → K2SO4 + Cr2(SO4)3 + 3Fe2(SO4)3+ 7H2O In ionic form the reaction can be written as Cr2O72- + 6Fe2+ + 14 H+ → 2 Cr3+ + 6Fe3++ 7H2O Orange

green

This redox reaction can be split apart in two parts- one showing the oxidation and the other reduction Cr2O72- + 6e¯ + 14H+ = 2 Cr3+ + 7H2O + VI

6 Fe2+ +II

+III

→ 6 Fe3+ +III

+ 6 e-

reduction (Oxidation number has decreased) oxidation (Oxidation number has increased)

The end point of the titration as indicated earlier has to be defined with the help of an indicator. Diphenylamine is one such indicator (internal indicator as it is added to the reaction mixture). The end point is marked with an intense blue violet colouration. The reduction potential value of this system is E°red = +0.76V which is very near to that of ferrousferric system (E°red = + 0.77V). Thus phosphoric acid is required in this titration as it reacts with the product ( yellow ) Fe3+ ion to form the complex ion [Fe(HPO4)]+, thus lowering the formal potential of the Fe(III)/Fe(II) system thereby increasing its reducing power. Thus the end point is sharper. The indicator action of diphenylamine can be understood as:

The diphenyl amine (I) undergoes oxidation first into a colourless diphenylbenzidine (II) which is the real indicator and is reversibly further oxidized to diphenylbenzidine violet (III). The titrations done here are double titrations. This means in the first titration K2Cr2O7 is standardized by titrating with a standard solution of Mohr’s salt and in the second titration this standardized K2Cr2O7 is titrated against a given oxalate ion solution of unknown strength. So experimental work involves three steps: Experiment no.1: AIM: To estimate the strength of given Mohr’s salt solution by titrating against Potassium dichromate solution. Standardize the Potassium dichromate solution with (approx. 0.025N) standard Mohr’s salt solution prepared by you. Step 1: Preparation of standard solution of Mohr’s salt (roughly N/40) Take a clean dry weighing bottle. Weigh it when empty with out the bottle lid. (as demonstrated in lab session) This is weight of empty weighing bottle (w1). Now put an approximate quantity of Mohr’s salt to be weighed in the weighing bottle. Weigh again. In case the added amount is too much than required , take out some with the help of a clean spatula and weigh the weighing bottle. In case the added amount is less than required, add some more into the weighing bottle. (Never add compound to the weighing bottle, when it is still placed on the pan of the weighing balance. Always take it out, and then add or remove the compound). Note down the weighing readings as weight of weighing bottle with the compound (w2) Transfer the compound into 100 mL measuring flask with the help of a funnel and keep the weighing bottle back on the pan of weighing balance. Note down the weight as weight of weighing bottle after transference of compound (w3). The difference of (w3) and (w2) gives the amount of compound actually transferred to the measuring flask. Dissolve the solid in minimum amount of water. Add a few drops of conc. H2SO4 acid. Shake and make up the volume till the 100 ml mark on the flask. Put the stopper on the flask and shake a number of times, vigorously to homogenize the solution Step 2: Standardisation of K2Cr2O7 soln. by titrating with standard Mohr’s salt solution Using Internal Indicator Diphenylamine • Pipette out 10 mL of Mohr’s salt solution in the conical flask . • Add approximately 10 mL of dil. H2SO4 to the same flask. • Add 2 drops of the indicator diphenylamine and 2 mL of 1:1 ortho H3PO4. • Titrate the reaction mixture with potassium dichromate solution taken in burette till a colour change from light green to blue- violet is obtained.



Repeat the titration for three concordant readings.

Using External Indicator Potassium Ferricyanide • Take a clean grooved tile. Add 2 drops of the indicator – Potassium Ferricyanide (2% aq. Soln.) in each groove. • Pipette out 10 mL of Mohr’s salt solution in the conical flask . • Add approximately 10 mL of dil. H2SO4 to the same flask. • Titrate the reaction mixture with potassium dichromate solution taken in burette. Stop at the point when some green colour in the solution starts appearing. • Dip a clean glass rod in the reaction mixture and then in the indicator taken on the grooved tile. • Notice the colour. A prussian blue (Turnbull’s blue) colouration or precipitate indicates the presence of reactant Fe2+ in the solution. • Continue the titration again with the same reaction mixture. Again after adding about 1 mL, repeat the indicator step and observe. • Repeat above two steps till no prussian blue colouration is observed when reaction mixture is added to the indicator. • Repeat the titration again, now checking the reaction mixture near end point after interval of 0.1 mL addition of Dichromate each time. • Repeat the titration for three concordant readings.

Step 3: Titration of Standardised K2Cr2O7 soln. with Mohr’s salt solution of unknown strength. • Repeat the above titration , now taking Mohr’s salt solution (whose concentration is to be determined) in place of standard Mohr’s salt solution.

Observations and Calculations Calculation of the amount of Mohr’s salt to be weighed for preparing the standard solution For preparing the standard solution of Mohr’s salt, first the amount required to weigh is calculated as Strength (g/L) = Normality x Equivalent weight. Normality of Mohr’s salt solution to be prepared is 0.025 N Equivalent weight of Mohr’s salt = Mol. Wt./ no. of equivalents = 392/ 1 = 392 Thus strength = 0.025 x 392 = 9.80 g/L For 100 mL solution, amount of Mohr’s salt to be weighed is 0.980 g.

OBSERVATION TABLE 1. Weighing observations: a. b. c.

Weight of empty weighing bottle; w1 (g) = ……………….. Weight of empty weighing bottle + Mohr’s salt; w2 (g) = ……………….. Weight of weighing bottle after transference of Mohr’s salt to the standard solution.; w3 (g) = ……………….. Amount of Mohr’s salt transferred to 100 ml measuring flask = ( w2 - w3) g Thus strength of Mohr’s salt solution is = ( w2 - w3) x 10 g/L And Normality of Mohr’s salt solution is = ( w2 - w3) x 10 N 392 (or Molarity Mohr’s salt solution is = ( w2 - w3) x 10 M) 392

2. Titration of K2Cr2O7 soln. Vs Mohr’s salt solution (self prepared) Volume of Mohr’s salt solution used in each titration = 10 mL Indicator used = Diphenylamine (2 drops) + 2ml 1:1 H3PO4 Colour change at end point = green to violet S. No. Initial

Burette readings R1 Final R2

Volume of K2Cr2O7 Consumed (R2 - R1) mL

Thus applying the normality relation NMohrVMohr = NdichrVdichr Thus Ndichr =

NMohrVMohr Vdichr ( Calculations alternatively can also be done using molarity concept 6MdichrVdichr = MMVM Mdichr

=

MMVM 6Vdichr )

3.

Titration of Standardised K2Cr2O7 soln. Vs Mohr’s salt solution (unkown conc.) Volume of Mohr’s salt solution used in each titration = 10 mL Indicator used = Diphenylamine (2 drops) + 2ml 1:1 H3PO4 Colour change at end point = green to violet

S. No. Initial

Burette readings R1 Final R2

Volume of K2Cr2O7 Consumed (R2 - R1) mL

Thus applying the normality relation NMohrVMohr = NdichrVdichr Thus NMohr =

NdichrVdichr VMohr ( Calculations alternatively can also be done using molarity concept 6MdichrVdichr = MMohrVMohr MMohr =

MdichrVdichr 6VMohr

)

Thus strength (g/L) of Mohr’s salt solution (of unknown conc.) = Normality x Eq. Wt Or

= Molarity x Mol.wt.

Result: The strength of given Mohr’s salt solution was found to be………………..g/L

Experiment no.2: AIM: To estimate the strength of given oxalic acid solution by titrating against Potassium permanganate solution. Standardize the Potassium dichromate solution with (approx. 0.025N) standard oxalic acid solution prepared by you. Step 1: Preparation of standard solution of oxalic acid (roughly N/40) Take a clean dry weighing bottle. Weigh it when empty with out the bottle lid. (as demonstrated in lab session) This is weight of empty weighing bottle (w1). Now put an approximate quantity of oxalic acid to be weighed in the weighing bottle. Weigh again. In case the added amount is too much than required , take out some with the help of a clean spatula and weigh the weighing bottle. In case the added amount is less than required, add some more into the weighing bottle. (Never add compound to the weighing bottle, when it is still placed on the pan of the weighing balance. Always take it out, and then add or remove the compound). Note down the weighing readings as weight of weighing bottle with the compound (w2) Transfer the compound into 100 mL measuring flask with the help of a funnel and keep the weighing bottle back on the pan of weighing balance. Note down the weight as weight of weighing bottle after transference of compound (w3). The difference of (w3) and (w2) gives the amount of compound actually transferred to the measuring flask. Step 2: Standardisation of KMnO4 soln. by titrating with standard oxalic acid solution • Pipette out 10 mL of oxalic acid solution in the conical flask . • Add approximately 10 mL of dil. H2SO4 to the same flask. • Heat the reaction mixture upto 60°C. You can have a rough idea about the temperature. This can be done by observing when the first bubble starts appearing in the solution and the flask is just unbearable to touch. • Titrate the hot reaction mixturequickly with potassium permanganate solution taken in burette till a first pink shade appears in the solution. • Repeat the titration for three concordant readings. Step 3: Titration of Standardised KMnO4 soln. with Mohr’s salt solution of unknown strength. • Repeat the above titration, now taking Mohr’s salt solution (whose concentration is to be determined) in place of standard Mohr’s salt solution.

Observations and Calculations Calculation of the amount of oxalic acid to be weighed for preparing the standard solution For preparing the standard solution of oxalic acid, first the amount required to weigh is calculated as

Strength (g/L) = Normality x Equivalent weight. Normality of oxalic acid solution to be prepared is 0.025 N Equivalent weight of oxalic acid = Mol. Wt./ no. of equivalents = 126/ 2 = 63 Thus strength = 0.025 x 63 = 1.56 g/L For 100 mL solution, amount of oxalic acid to be weighed is 0.0156 g. OBSERVATION TABLE 2. Weighing observations: a. b. c.

Weight of empty weighing bottle; w1 (g) = ……………….. Weight of empty weighing bottle + oxalic acid; w2 (g) = ……………….. Weight of weighing bottle after transference of oxalic acid to the standard solution.; w3 (g) = ……………….. Amount of oxalic acid transferred to 100 ml measuring flask = ( w2 - w3) g Thus strength of oxalic acid solution is = ( w2 - w3) x 10 g/L And Normality of Mohr’s salt solution is = ( w2 - w3) x 10 N 63 (or Molarity Mohr’s salt solution is = ( w2 - w3) x 10 M) 126

2. Titration of KMnO4 soln. Vs oxalic acid solution (self prepared) Volume of oxalic acid solution used in each titration = 10 mL Indicator used = KMnO4 as self indicator Reaction temperature = ~60°C Colour change at end point = colourless to pink S. No. Initial

Burette readings R1 Final R2

Thus applying the normality relation NoxVox = NpermangVpermang

Volume of KMnO4 Consumed (R2 - R1) mL

Thus Npermang =

NoxVox Vpermang ( Calculations alternatively can also be done using molarity concept 5MpermangVpermangr = 2MoxVox Mpermang

3.

=

2MoxVox 5Vpermang )

Titration of Standardised KMnO4 soln. Vs oxalic acid solution (unkown conc.) Volume of oxalic acid solution used in each titration = 10 mL Indicator used = KMnO4 as self indicator Reaction temperature = ~60°C Colour change at end point = colourless to pink

S. No. Initial

Burette readings R1 Final R2

Volume of KMnO4 Consumed (R2 – R1) mL

Thus applying the normality relation NoxVox = NpermangVpermang Thus Nox =

Npermang Vpermang Vox ( Calculations alternatively can also be done using molarity concept 5MpermangVpermang = 2MoxrVox Mox =

2 MpermangVpermang 5Vox )

Thus strength (g/L) of Oxalic acid solution (of unknown conc.) = Normality x Eq. Wt Or

= Molarity x Mol.wt.

Result: The strength of given Oxalic acid solution was found to be………………..g/L