Download - Expt 12- 19

LECTURE NOTES: Reduction- Oxidation Titration

(Expt. 12- 19)

1. Neutralization- acid and base2. Redox- reducing agent, oxidizing agent

Oxidizing agents: KMnO4 ; ceric sulfate; Iodine ; bromine; potassium iodate

Reducing agents: Oxalic acid; ferrous sulfate; sodium thiosulfate

Note:

h= no. of electrons gained or lost depending on its role in redox

EXPERIMENT 12:Preparation and Standardization of Potassium Permanganate Solution

EXPERIMENT 12:A. Preparation of Potassium Permanganate

Solution/ Class

1. Weigh roughly 5g of KMnO4

2. Dissolve it in 1500 mL distilled water in a big beaker.

3. Cover with a watch glass and boil the solution for 15 mins

(prevent evaporation and accelerate destruction of organic matters or decomposition)

4. Allow to stand for at least 2 days (allow decomposition of organic matter to proceed to completion)

5. Filter through asbestos into an amber color bottle(get rid of decomposition inorganic product which is MnO2, a brown ppt. MnO2

catalyzes the auto decomposition of KMnO4)

EXPERIMENT 12:LECTURE NOTES:

Asbestos- carcinogenic for filtration of inorganic matter should be lined with a glass wool as support

(funnel) when used as a filter.

Should be washed with distilled water until water coming out of the funnel is clear before KMnO4 is filtered.

EXPERIMENT 12:B. Standardization of Potassium Permanganate

Solution

1. Weigh 0.1-0.2g of Na2C2O4 (weighing bottle)

2. Dissolve it in 250 mL distilled water in a big beaker.

3. Add 7mL of conc. sulfuric acid

4. Heat the solution to 70 C ⁰ (maintain ---reaction between sodium oxalate and KMnO4 occurs stoichiometrically in a hot acidic medium)

5. Titrate with KMnO4 solution (Geissler) until a light pink color persists for 15 seconds is obtained. ( Geissler: IR, FR)

6. One determination per pair

7. Compute for Mean N and M as a class

EXPERIMENT 12:LECTURE NOTES

Standardization: primary standardization

Type of Titration: Redox

Method of Titration: Direct Permanganate method

1. Na2C2O4 - primary standard, reducing agent

2. KMnO4 -strong oxidizing agent

-reacts rapidly with reducing substance in

acidic medium

- serves as indicator itself due to its intense

color imparting a pink coloration

-disadvantage: unstable

EXPERIMENT 12:Lecture Notes

3. H2SO4 -to keep the acid conc. constant preventing the formation of MnO2

- supply hydrogen ion

4. Hot acidic medium- the reaction between Na2C2O4 and KMnO4 stoichoimetrically occur

in a hot acidic medium.


5. Equation:Molecular

lost=1eX 2= 2e x5= 10e +3 +4

2KMnO4 + 5Na2C2O4 + 8H2SO4 2MnSO4 + 5Na2SO4 + 10CO2 + K2SO4 + 8H2O

+7 gained=5e x 2 = 10e +2

Half ionic equation

C2O4 -2 CO2

MnO4- Mn+2

1. Balance the atoms

2. Balance the charges

3. Balance the number of electrons gained/lost


6. Computation: Normalitymeq KMnO4 = meq Na2C2O4

(NxmL) KMnO4 = (g /MEW) Na2C2O4

N KMnO4 = g Na2C2O4

MW

h X 1000

mL KMnO4

NOTE:

h= no. of electrons gained or lost depending on its role in redox

KMnO4 (oxidizing agent)= 5 electrons gained Na2C2O4 (reducing agent)= 2 electrons lost


6. Computation: Molarity

mn KMnO4 = mn Na2C2O4 x rr

(MxmL) KMnO4 = (mg /MW) Na2C2O4 X 2KMnO4

5Na2C2O4

M KMnO4 = mg/MW Na2C2O4 X 2KMnO4

5Na2C2O4

mL KMnO4

EXPERIMENT 13:Preparation and Standardization of

Oxalic acid Solution

EXPERIMENT 13:A. Preparation of Oxalic Acid Solution/ Class-

Group 5

1. Weigh roughly 9.6g of oxalic acid

2. Dissolve it in distilled water to make 1500 mL

3. Store in an amber colored bottle

EXPERIMENT 13:B. Standardization of Oxalic Acid Solution

1. Run down 30 mL of oxalic acid (Mohr) into a big beaker (Mohr : IR, FR)

2. Dilute with 200 mL distilled

3. Add 7mL of conc. sulfuric acid

4. Heat the solution to 70 C ⁰ (maintain---reaction between oxalic acid and KMnO4 occurs stoichiometrically in a hot acidic medium)

5. Titrate with KMnO4 solution (Geissler) until a light pink color persists for 15 seconds is obtained. (Geissler : IR, FR)


7. Compute for Mean N and M as class


Standardization: Secondary standardizationType of Titration: RedoxMethod of Titration: Direct Permanganate method

1. KMnO4 – secondary standard

-strong oxidizing agent-reacts rapidly with reducing substance (oxalic acid) in hot acidic medium- serves as indicator itself due to its intense

color imparting a pink coloration -disadvantage: unstable



lost= 1ex2 X 5= 10e

+3 +42KMnO4 + 5H2C2O4 + 3H2SO4 2MnSO4 + 10CO2 + K2SO4 + 8H2O

+7 gained = 5ex2= 10e +2

Half ionic equation

C2O4 -2 CO2

MnO4- Mn+2

1. Balance the atoms2. Balance the charges3. Balance the number of electrons gained/lost


3. Computation: Normality

meq H2C2O4 = meq KMnO4

(NxmL) H2C2O4 = (NxmL) KMnO4

N H2C2O4 = N x mL KMnO4

mL H2C2O4



mn H2C2O4 = mn KMnO4 x rr

(MxmL)H2C2O4 = (MxmL) KMnO4 X 5H2C2O4

2KMnO4

M H2C2O4 = MxmL KMnO4 X 5H2C2O4

2KMnO4

mL H2C2O4

EXPERIMENT 14:Assay of Sodium Nitrite


1. Weigh 0.4 to 0.5g NaNO2 accurately

2. Dissolve it in enough distilled water to make 100mL in a volumetric flask

3. Pipet 10 mL NaNO2 solution (wash pipet first)

4. Introduce it into a beaker containing a mixture of:

25 mL KMnO4 (Geissler buret: IR, FR)

100 mL distilled water

5 mL conc. H2SO4

by immersing the tip of the pipet beneath the surface of the permanganate mixture [prevent the escape of nitrous acid (HNO2)]


5. Warm the liquid to 40⁰C and allow it to stand for 5 minutes. (to accelerate the oxidation of HNO2 to HNO3)

6. Add 25 mL of standard oxalic acid solution (Mohr: IR, FR)- (KMnO4 is decolorized because the oxalic acid was added in excess)

7. Heat to 80⁰C8. Titrate with standard KMnO4 to a light pink end

point ( Geissler: IR,FR)

9. One determination /pair

10. Look for USP specifications for NaNO2

11. Compute Mean % purity as a class

12.Disposition


Assay of Sodium Nitrite


Method of Titration: Residual Permanganate method

1. NaNO2- reducing agent

2. KMnO4 – oxidizing agent

Note : Ferrous, Calcium, Lead (II) and titanium compounds are also assayed by permanganate method



2NaNO2 + H2SO4 2HNO2 + Na2SO4

lost= 2e x5= 10e

+3 +5

5HNO2 +2KMnO4+ 3H2SO4 5HNO3+ 2MnSO4+ K2SO4+ 3H2O

+7 +2 gained= 5e X2=10e

lost= 1ex2 X 5= 10e

+3 +42KMnO4 + 5H2C2O4 + 3H2SO4 → 2MnSO4 + 10CO2 + K2SO4 + 8H2O +7 gained = 5ex2= 10e +2


2. Equation:Half ionic equation

NO2- NO3

-

MnO4- Mn+2

C2O4 -2 → CO2

MnO4- → Mn+2

EXPERIMENT 14: LECTURE NOTES


meq NaNO2 = meq KMnO4 that reacts with NaNO2

gpure NaNO2= [(NxmL) Total KMnO4 - (N x 25mL) H2C2O4] MW

hx1000

g pure NaNO2= [(N x mL) Total KMnO4 -(N x 25mL) H2C2O4] x MW

h x 1000

% NaNO2(w/w) = g pure NaNO2 _______X 100

0.4 to 0.5g = X (aliqout)

100 mL 10 mL



mn NaNO2 = mn KMnO4 that reacts with NaNO2 x rr

gNaNO2= [(MxmL) Total KMnO4 - (M x 25mL) H2C2O4 X KMnO4 ] X HNO2 X NaNO2 X MW

H2C2O4 KMnO4 HNO2 1000

gNaNO2= [(M x mL) Total KMnO4 -(M x 25mL) H2C2O4 X 2 KMnO4 ] x 5HNO2 X 2NaNO2 x MW

5 H2C2O4 2KMnO4 2HNO2 1000

% NaNO2(w/w) = g pure NaNO2 X 100

0.4 to 0.5g = X (aliqout)

100 mL 10 mL

\


Sodium thiosulfate Solution


Experiment 15-18:


Method: Iodimetry/ Iodometry1. Iodimetry- is the process wherein a standard solution of iodine is the

titrating agent and it acts as an oxidizing agent.

Pharmaceutical solutions assayed Iodimetrically:

2. Arsenites 6. ascorbic acid

3. Sb+3 compds 7. methenamine

4. Thiosulfates

5. Sulfites

6. Mercurous compds


2. Iodometry- is a process where in the sample of an oxidizing agent is made to liberate an equivalent amount of iodine from KI which is titrated with Na2S2O3 solution.

Pharmaceutical solutions assayed Iodometrically:

1. Iron 6. arsenous

2. Copper 7. chlorine

3. Manganese 8. Bromine

4. Chromium 9. Iodine

5. Cobalt

EXPERIMENT 15:A. Preparation of Starch Solution

1. Triturate 1g of arrowroot starch with 10 mL of distilled water. (starch paste)

2. Boil 200 mL of distilled water

3. Add the starch paste to it with constant stirring

4. Boil the mixture gently until it forms a thin, translucent liquid.


Starch test solution- serves a indicator

Starch grains contain:

1. α- amylose (water insoluble)- with Iodine forms a violet color

2. ß-amylose (water soluble)- with Iodine forms blue color


1. Larger grains- arrow root and potatoes( will give more ß- amylose)

2. Smaller grains- rice and corn ( will give lesser ß- amylose)

EXPERIMENT 15:B. Preparation of Sodium Thiosulfate Solution/

Class

1. Weigh roughly 26 g of sodium thiosulfate and 0.4 g of sodium carbonate (preservative to prevent acid catalyzed hydrolysis)

2. Dissolve this in 1000 mL of recently boiled and cooled distilled water (for bacterial sterilization and to expel CO2)

3. Store in an amber colored bottle and label.

EXPERIMENT 15:C. Standardization of Sodium Thiosulfate

Solution

1. Weight accurately 1.1g of primary standard KIO3 into a 500-mL volumetric flask;

2. Dissolve in about 200 mL of distilled water. Dilute to the mark and mix thoroughly.- (one per class)

3. Pipet 50.0 mL aliquot of standard KIO3 solution into a 250 mL conical flask ( one per paired groups)

4. Introduce 2g of KI (iodate free) and swirl the flask to speed up solution.


Solution

5. Add 2 mL of 6M HCl (do not add if not Na2S2O3 is not yet ready)- catalyze the liberation of I2

6. Immediately titrate with sodium thiosulfate until solution is pale yellow (reduction of the amount of Iodine)

7. Introduce 5 mL of the starch solution

(There will be a formation of big lumps of the blue iodo starch complex if the indicator is added before the titration with Na2S2O3 to the pale yellow endpoint. This will result to difficulty of the blue color to disappear.)


Solution

8. Continue the titration with standard sodium thiosulfate to the disappearance of the blue color.


10. Compute the Mean N and M of Na2S2O3 as a class


Standardization: primary standardization


Method of Titration: Iodometry

1. KIO3- primary standard (oxidizing agent)

2. Na2S2O3 – titrant; reducing agent


3. Equations:Molecular:

+5 gained=5eX2= 10e 0

• 2KIO3+ 10 KI + 12HCl 6I2 + 12KCl + 6H2O

-1 lost= 1 x2=2eX 5= 5e 0

reduced to lowest term:

+5 gained = 5e x2= 10e 0

• KIO3 + 5KI + 6HCl → 3I2 + 6KCl + 3H2O

-1 0

lost= 1 x 2=2e X5= 10e



+5 gained=5eX2= 10e 0

• KIO3+ 5 KI + 6HCl 3I2 + 6KCl + 3H2O

-1 lost= 1 x2=2eX 5= 5e 0

0 gained = 1x2=2e X 1= 2e -1

• I2 + 2Na2S2O3 → 2NaI +Na2S4O6

+2 +2.5

lost= 0.5 x 2=1e X2= 2e


3. Equations:Half- ionic equation:

• IO3- → I2

I- → I2

• S2O3-2 → S4O6

-2

I2 → I-


3. Equations:Half- ionic equation:

• IO3- + 5 I-→ 3I2+ 3H2O

• 2S2O3-2 + I2

→ S4O6-2 + 2I-


4. Computation: Normalitymeq Na2S2O3 = meq KIO3

(NxmL) Na2S2O3 = (g /MEW) KIO3

N Na2S2O3 = g KIO3

MW

h X 1000

mL Na2S2O3

Note: g= 1.1g = x

500 mL 50mL h=6



mn Na2S2O3 = mn KIO3 x rr

(MxmL) Na2S2O3 = mg KIO3 x 3 I2 x 2Na2S2O3

MW 1 KIO3 1 I2

M Na2S2O3 = mg KIO3 X 3 I2______ x 2Na2S2O3

MW___ 1 KIO3 1 I2

mL Na2S2O3


Iodine Solution

EXPERIMENT 15:A. Preparation of Iodine Solution/Class

1. Weigh roughly 14.0g of Iodine crystals

2. Dissolve it in a solution of 36 g KI in 400 mL of distilled water. (KI is a solubilizing agent to increase the solubility of Iodine crystals)

3. Add 12 drops of 6N HCl (neutralized any alkali present in KI)

4. Add enough distilled water to complete the volume to 1000 mL

5. Store in an amber colored bottle.

EXPERIMENT 16:B. Standardization of Iodine Solution

1. Run down 30 mL of sodium thiosulfate (Mohr: IR, FR)

2. Dilute with 100 mL distilled water

3. Add 5 mL of starch solution

4. Titrate with standard iodine solution to a blue color endpoint. (Geissler: IR and FR)


6. Compute Mean N and M as a class.


Standardization: secondary standardization


Method of Titration: Iodimetry

1. Na2S2O3 solution - secondary standard; reducing agent

2. Iodine – titrant; oxidizing agent



+2 lost=0.5 x2= 1e x 2= 2e + 2.5

2Na2S2O3+ I2 2NaI+ Na2S4O6

0 -1

gained= -1x2= 2e x1= 2e


3. Equations:Half ionic :

• S2O3-2 S4O6

-2

I2 I-1

2S2O3-2 + I2

→ S4O6-2 + 2I-



meq I2 = meq Na2S2O3

(NxmL) I2 = (NxmL) Na2S2O3

N I2 = N x mL Na2S2O3

mL I2



mn I2 = mn Na2S2O3 x rr

(MxmL) I2 = (MxmL) Na2S2O3 X 1 I2 ___ 2

Na2S2O3

M I2 = MxmL Na2S2O3 X 1 I2 ___

2 Na2S2O3

mL I2

EXPERIMENT 17:Assay of Tartar Emetic

EXPERIMENT 17:Assay of Tartar Emetic

1. Weigh 0.4 to 0.5g tartar emetic accurately (tared flask)

2. Dissolve it in 30 mL distilled water

3. Add 25 mL of saturated solution of sodium bicarbonate (neutralize the HI formed thus preventing a reversible reaction and allowing the reaction to proceed to completion)

4. Add 5 mL of starch T.S.

5. Titrate with standard iodine solution to a blue color end point. (Geissler: IR and FR)


7. Compute for the mean % purity of tartar emetic as a class.

8. Look for USP/ NF specs

9. Disposition


1. Tartar Emetic (KOSbC4H4O6. ½ H2O)

• double salt• Reducing agent

2. Method of Titration: Direct Iodimetric method


+3 lost= 2e- +5

KOSbC4H4O6 + I2 + 2NaHCO3 KO2SbC4H4O6 + 2NaI +2CO2 +H2O

0 gained= 1x2= 2e -1


3. Equation:Half ionic

Sb +3 Sb +5

I2 I-1



meq KOSbC4H4O6. ½ H2O = meq I2

(g /MEW) KOSbC4H4O6. ½ H2O = (NxmL) I2

g pure KOSbC4H4O6. ½ H2O = (NxmL) I2 X MW

h X 1000

% KOSbC4H4O6. ½ H2O= g pure KOSbC4H4O6. ½ H2O x 100

g impure or g sample



mn KOSbC4H4O6. ½ H2O = mn I2 x rr

g pure KOSbC4H4O6. ½ H2O = (MxmL) I2 x 1 KOSbC4H4O6. ½ H2O

MW 1 I2

1000

gpure KOSbC4H4O6. ½ H2O = (MxmL) I2 x 1 KOSbC4H4O6. ½ H2O X MW

1 I2 1000

% KOSbC4H4O6. ½ H2O (w/w) = g pure KOSbC4H4O6. ½ H2O X 100


EXPERIMENT 18:Assay of Cupric Sulfate


1. Weigh 0.4 to 0.5g cupric sulfate accurately (tared iodine flask)

2. Dissolve it in 50 mL distilled water

3. Add 4 mL of 6 N acetic acid

4. Add 3 g of KI (Note the color change after adding: reddish brown color indicates plenty of I2 is liberated)

5. Titrate the liberated iodine with sodium thiosulfate until a golden yellow color is obtained. ( Mohr: IR) (indicates the reduction of the amount of I2)

6. Add 3 mL of starch T.S. (There will be a formation of big lumps of the blue iodo starch complex if the indicator is added before the titration with Na2S2O3 to the golden yellow endpoint. This will result to difficulty of the blue color to disappear.)


7. Continue the titration with std. sodium thiosulfate solution until the disappearance of blue color. (FR)


9. Compute for the % purity of cupric sulfate as a class.

10. Look for USP/NF specs

11. Disposition


Assay of Cupric sulfate

1. CuSO4.5H2O (oxidizing agent)

• blue vitriol

2. KI (reducing agent)

3. Type of Titration: Redox

4. Method of Titration: Iodometry


5. Equations:Molecular

+2 gained= 1e X2 =2e +1

2CuSO4.5H2O +4KI 2CuI + I2 + 2K2SO4 +10 H2O

lost= 1e x2=2e X1= 2e

0 gained= 1ex2=2e x1= 2e -1

I2 + 2Na2S2O3 2NaI + Na2S4O6

+2 lost= 0.5 X2= 1e X2=2e +2.5


5. Equation:Half ionic

Cu +2 Cu +1

I2 I-1



meq CuSO4O6. 5 H2O = meq Na2S2O3

(g /MEW) CuSO4. 5H2O = (NxmL) Na2S2O3

g pure CuSO4 5 H2O = (NxmL) Na2S2O3 X MW CuSO4 5 H2O

h X1000

% CuSO4. 5H2O= g pure CuSO4. 5H2O x 100




mn CuSO4. 5H2O = mn Na2S2O3 x rr

g pure CuSO4. 5H2O = (MxmL) Na2S2O3 x 1 I2 __________ X 2 CuSO4. 5H2O

MW 2 Na2S2O3 1 I2

1000

gpure CuSO4. 5H2O = (MxmL)Na2S2O3 x 1 I2 __________ X 2 CuSO4. 5H2O X MW

2 Na2S2O3 1 I2 1000

% CuSO4. 5 H2O (w/w) = g pure CuSO4. 5H2O X 100


EXPERIMENT 19:ASSAY OF ASCORBIC ACID

EXPERIMENT 19:A. Preparation of Standard Potassium

Bromate/Class- Group 8

1. Weigh approximately 1.6 g into a 1000 mL volumetric flask.

2. Dissolve the KBrO3 in about 400-mL of distilled water.

3. Dilute to the mark, mix thoroughly.

4. Keep in an amber bottle.

EXPERIMENT 19:Assay of Ascorbic Acid

1. Weigh accurately 3 to 5 vitamins

2. Pulverize them thoroughly in a mortar, and transfer the powder to a dry weighing bottle.

3. Weigh accurately 0.4 to 0.5 g sample into a dry 250 mL conical flask (with cover)

4. Dissolve the sample in 50mL of 1.5 M H2SO4 (freshly

prepared; converts BrO3 to Br2); then add about 5g of KBr. ( will produce excess Br2)

5. Titrate immediately with standard KBrO3 to the first faint yellow due to excess Br2 (Geissler: IR, FR)

EXPERIMENT 19:Assay of Ascorbic Acid

6. Record the volume of KBrO3 used.

7. Add 3g of KI and 5mL of starch indicator; back titrate with standard Na2S2O3 to the disappearance of blue color. (titration should be done without delay to prevent the air oxidation of ascorbic acid)

8. Calculate the ave. mass (in mg) of ascorbic acid tablet.

9. Two determination /pair

10. Compute Mean % purity per pair

11. Look for USP specifications for ascorbic chewable tablets

12. Disposition

EXPERIMENT 19:LECTURE NOTES:

Assay of Ascorbic Acid

1.Type of Titration: Redox

2. Method: Bromination

Applicable for chewable vitamin C (not coated tablets)

The binder in most Vitamin C tablets remains in the suspension through out the analysis. If the binder is starch, the characteristic color of the complex with I2 appears upon the addition of KI

The volume of Na2S2O3 needed for the back titration seldom exceeds a few millimeters.


3. Equation:

BrO3 + 5Br- + 6H+→ 3Br2 + 3H2O

4. + Br2 → + 2 Br- + 2H+

Ascorbic acid (C6H8O6) dehydro ascorbic acid (C6H6O6)

Br2 + 2 I-→ Br- + I2

I2 + 2 S2O3-2 → 2 I- + S4O6

-2



meq AA = meq Br2 that reacts with AA

g/MEWAA= g/MEW KBrO3 - g/MEW Na2S2O3

= whole Br2 - x’ss Br2

gpure AA = [ (NxmL) KBrO3 - (N x mL) Na2 S2O3 ] MW

hx1000

g pure AA= [ (N x mL)KBrO3 - (N x mL) Na2S2O3] x MW AA

h x 1000

g pure AA___________ = X________________ g sample (0.4 to 0.5 g) ave. wt of 3- 5 tablets

x= amount (g) of AA/ tablet ; convert to mg/ tablet


Computation: Normality KBrO3

NKBrO3 = 2.6g KBrO3

MW

h X 1000

1000 mL

Note: h=6

Equation:

BrO3 → Br2



mn AA = mn Br2 that reacts with AA x rr

mn AA= mn KBrO3 - mn Na2S2O3 x rr

= whole Br2 - x’ss Br2 x rr

gpure AA = [(MxmL) KBrO3 x 3Br2 - (M x mL) Na2 S2O3 X 1 I2______ x 1Br2 ] MW 1KBrO3 2Na2S2O3 1I2

1000

g pure AA= [(M x mL)KBrO3 x 3Br2 - (M x mL) Na2S2O3 X 1 I2________ X 1Br2] x MW AA

1KBrO3 2 Na2S2O3 1 I2 1000

g pure AA___________ = X________________ g sample (0.4 to 0.5 g) ave. wt of 3- 5 tablets

x= amount (g) of AA/ tablet ; convert to mg/ tablet


Computation: Molarity KBrO3

MKBrO3 = 2.6g KBrO3

MW

1000

1L

Download - Expt 12- 19

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