This document describes the process of estimating the amount of vitamin C in a sample using iodine titration. Key steps include: 1. Preparing standard solutions of potassium dichromate, sodium thiosulfate, and iodine through calculations of amounts needed. 2. Standardizing the iodine solution against the sodium thiosulfate solution. 3. Titrating a 0.125g sample of vitamin C that has been dissolved in metaphosphoric acid with the standard iodine solution. 4. Calculating the amount of vitamin C in the sample based on the titration volume and normality of iodine. The results found 101mg of vitamin C in the

1. Presented By
Sajjad Alam
Department of Chemistry
University of Rajshahi
ID : 1910323114
Estimation of the amount of
Vitamin C (L-Ascorobic acid)
in the supplied sample.

2. 2
Introduction
 Scientific Name: Ascorbic Acid
 Chemical Formula: C6H8O6
 Appearance: White to Slightly Yellowish Crystalline powder; practically
odorless with a strong acidic taste.
 Melting Point: About 1900c
 Boiling Point: 5530c
 Solubility in water: 330 g/L; Dissolves well in water to give mildly acidic solution.
 One of the most well-known properties of vitamin C is its antioxidant activity.
 Iron Absorption: Another notable property of vitamin C is its ability to enhance
the absorption of non-heme iron from plant-based sources.
 Heat and Light Sensitive.
 Growth and Repair tissue in all parts of the body.
 Supports immune fraction and protect from viral diseases.
 Prevent from getting scurvy.

3. Objective
3
The primary objective of this method is to
determine the quantitative amount of vitamin C
(ascorbic acid) present in a sample.
Quantitative determination
The method aims to specifically measure the
content of vitamin C without interference from
other compounds or substances present in the
sample. This specificity ensures that the results
obtained reflect the true amount of vitamin C
present.
Specificity
The method aims to provide accurate and precise
results. By carefully controlling the experimental
conditions and following standardized
procedures, the method aims to minimize errors
and variability in the results.
Accuracy and precision
By employing the same method and standardized
procedures, it becomes possible to assess and
compare the vitamin C content of various
samples, such as different food products or
dietary supplements.
Comparability
The method can be used for quality control
purposes, particularly in industries involved in the
production of vitamin C-rich products or dietary
supplements. It helps ensure that the products
meet regulatory standards or internal quality
specifications.
Quality control
The iodine titration method is often employed for
routine analysis, where a large number of
samples need to be analyzed within a short
period. Its simplicity, cost-effectiveness, and
relatively fast analysis time make it suitable for
routine testing and monitoring of vitamin C
content
Routine analysis

4. Application of Vitamin C Estimation
4
Food Industry
Pharmaceuticals
Clinical Research
Nutritional Research
Biomedical Research
Nutraceuticals
Quality Control
Regulatory Compliance

5. 5
Factors Affecting Vitamin C Estimation
Sample Preparation
Improper sample preparation can lead to
inaccurate results. Factors such as
inadequate homogenization, incomplete
extraction of vitamin C from the sample
matrix, or degradation of vitamin C
during sample preparation can affect the
analysis. It is crucial to follow appropriate
sample preparation techniques to ensure
representative and reliable results.
Temperature
Temperature: Temperature can impact
the stability of vitamin C. Higher
temperatures can accelerate the
degradation of vitamin C, leading to
underestimated results. Therefore,
samples should be stored and analyzed
under controlled temperature conditions
to preserve the integrity of vitamin C.
The reaction between vitamin C and
iodine is pH-dependent. Deviations from
the recommended pH range (usually
around pH 3) can affect the rate of
reaction and the stability of the vitamin
C-iodine complex. pH control, typically
achieved using buffer solutions, is
important for accurate analysis.
pH
Vitamin C is susceptible to oxidation,
particularly in the presence of oxygen,
light, and heat. Exposure to these factors
can lead to the degradation of vitamin C,
resulting in lower measured values. It is
important to handle samples and
reagents in a way that minimizes
oxidation during the analysis.
Oxidation

6. How Can We Minimize Errors
6
To minimize errors :
 Using fresh reagents,
 Precise pipetting and measurement.
 Proper indicator usage.
 Control pH .
 Environmental controls.
 Keeping equipment clean and calibrated.
 Performing multiple trials to ensure consistent results.
 Documentation and record-keeping

7. Preparation
Preparation of 0.1 N K2Cr2O7 Solution :
Calculation of the amount of K2Cr2O7 needed
Molecular weight of K2Cr2O7 = 294.18
Equivalent weight of K2Cr2O7 =
294.18
6
= 49.03 g/eq
The amount of K2Cr2O7 was required ,
W=
𝑆×𝑀×𝑉
1000
=
0.1×49.03×250
1000
= 1.225 g
Amount of K2Cr2O7 was taken = 1.224 g
7

8. Preparation
Preparation of 0.1 N K2Cr2O7 Solution :
Table: Data for the preparation of 250 mL 0.1 N K2Cr2O7 Solution
Calculation:
∴ Actual conc. Of K2Cr2O7 =
𝑇𝑎𝑘𝑒𝑛 𝑚𝑎𝑠𝑠
𝑟𝑒𝑞𝑢𝑖𝑟𝑒𝑑 𝑚𝑎𝑠𝑠
× 0.1 𝑁 =
1.224 𝑔
1.225 𝑔
× 0.1 𝑁 = 0.09 N
Required Mass
(g)
Mass Taken
(g)
Actual concentration
Of K2Cr2O7 (N)
1.225 1.224 0.09
8

9. Preparation
Preparation of 0.1 N Na2S2O3 Solution :
Calculation of the amount of Na2S2O3 needed
Molecular weight of Na2S2O3 = 248
Equivalent weight of Na2S2O3 =
248
1
= 248 g/eq
The amount of Na2S2O3 was required ,
W=
𝑆×𝑀×𝑉
1000
=
0.1×248×250
1000
= 6.20 g
Amount of K2Cr2O7 was taken = 6.22 g
9

10. Preparation
Preparation of 0.1 N Na2S2O3 Solution :
Table: Data for the preparation of 250 mL 0.1 N Na2S2O3 Solution
Calculation:
∴ Actual conc. Of Na2S2O3 =
𝑇𝑎𝑘𝑒𝑛 𝑚𝑎𝑠𝑠
𝑟𝑒𝑞𝑢𝑖𝑟𝑒𝑑 𝑚𝑎𝑠𝑠
× 0.1 𝑁 =
6.22 𝑔
6.20 𝑔
× 0.1 𝑁 = 0.10 𝑁
Required Mass
(g)
Mass Taken
(g)
Actual
concentration
Of Na2S2O3 (N)
6.20 6.22 0.10
10

11. Standardization
Standardization of the approximate 0.10 N Na2S2O3 solution with Standard 0.09 N K2Cr2O7 solution
 50 mL water + 1 g NaHCO3 + 2-3 mL conc. HCl in a conical flask.
 Add 1.5 g KI solution.
 Add 10 mL K2Cr2O7 solution with a pipette, cover with a watch glass and keep in dark for 5 minutes.
 Wash down the watch glass and side of flask with water.
 Titrate with Na2S2O3 from a burette till the red color fades.
 Add 1-2 mL starch indicator.
 Continue titration with Na2S2O3 solution till the color becomes green.
End point: Blue→ Green.
Cr2𝑂7
2−
+ 14H+ +6I → 2Cr3+ 7H2O + 3I2
6S2O3
2- + 3I2 → 3S4O6
2- + 6I-
11

12. Standardization
Standardization of the approximate 0.1 N Na2S2O3 solution with Standard 0.1 N K2Cr2O7 solution
Trial
No
Volume of
K2Cr2O7
/mL
Burette Reading of
Na2S2O3 /mL Difference
/mL
Average
Volume of
Na2S2O3 /mL
Concentration
of Na2S2O3
/N
Initial Final
1
10
0.40 11.35 10.95
10.92 0.09
2 11.35 22.35 11.00
3 22.35 33.15 10.80
Table: Standardization of Na2S2O3 iodometrically with standard K2Cr2O7 solution.
Calculation :
N1×V1=N2×V2
N1=
𝑁2×𝑉2
𝑉1
=
0.09×10
10.92
=0.08 N
∴ Therefore the strength of the supplied Na2S2O3 was 0.08 N 12

13. Preparation
Preparation of 0.1 N I2 Solution :
Calculation of the amount of I2 needed
Molecular weight of I2 = 254
Equivalent weight of I2 =
254
2
= 127 g/eq
The amount of I2 was required ,
W=
𝑆×𝑀×𝑉
1000
=
0.1×127×1000
1000
= 1.27 g
13

14. Preparation
Iodine solutions are prepared using potassium iodide​ because of the following reasons:
 Potassium Iodide (KI) converts the insoluble iodine( I₂) to soluble Iodide ion( I3
-).
 Concentration of iodine is increases in the solution.
 Lugol's solution:
KI mixed with I2 forms Lugol's solution.
Lugol's solution is an equimolar solution of iodine and potassium iodide.
Reaction :
KI + I₂ → KI₃
I2 insoluble in water.
KI3 contains I3
- ion, which is soluble in water.
14

15. Standardization
Standardization of the approximate 0.1 N I2 solution with Standard
0.1 N Na2S2O3 solution:
 Using a pipette filler, fill the 10 mL pipette with the iodine solution, and
transfer the contents of the pipette to the conical flask and add about
20 mL of water.
 Add a few drops of the starch indicator solution just prior to the end-
point, when the colour of the solution fades to pale yellow. Upon
addition of the indicator a blue-black colour should appear. The
thiosulfate solution should now be added dropwise, with thorough
swirling.
 The end-point of the titration is detected by a colour change from blue-
black to colourless. Note the burette reading.
2S2O3
2- + I2 → 2I- + S4O6
2-
15
Iodine Solution

16. Standardization
Trial No
Volume of I2
/mL
Burette Reading of
Na2S2O3 /mL Difference
/mL
Average Volume
of Na2S2O3 /mL
Concentration
of I2 /N
Initial Final
1
10
2.90 15.70 12.80
12.77 0.11
2 15.80 28.50 12.70
3 28.60 41.40 12.80
Table: Standardization of I2 iodometrically with standard Na2S2O3 solution.
Calculation :
N1×V1=N2×V2
N1=
𝑁2×𝑉2
𝑉1
= 0.08×12.77
10
=0.10 N
∴ Therefore the strength of the supplied I2 was 0.11 N
The normality factor of 0.1 N I2 Solution is, F=
𝐴𝑐𝑡𝑢𝑎𝑙 𝑆𝑡𝑟𝑒𝑛𝑔𝑡ℎ
𝑇ℎ𝑒𝑜𝑟𝑖𝑖𝑐𝑎𝑙 𝑆𝑡𝑟𝑒𝑛𝑔𝑡ℎ
=
0.10 𝑁
0.1 𝑁
= 1 16
Iodine Solution

17. Estimation of Vitamin C
Titration of the supplied Vitamin C with Standard 0.1 N I2 Solution:
 0.125 g Sample+ 200 mL metaphosphoric acid + 5 mL starch indicator.
 While mixing starch the solution should be kept on constant stirring.
 Finally titrate with iodine solution from burette.
 Color of end point should be bluish green.
• The chemical structure and antioxidant (reducing) action of
ascorbic acid are illustrated in the redox half equation above.
17

18. Final Titration
Titration of the supplied Vitamin C with Standard 0.1 N I2 Solution:
Trial
No
Amount of
Vitamin C
/g
Burette Reading
of I2 /mL Difference
/mL
Initial Final
1 0.125 0.72 12.19 11.47
 Ascorbic acid is susceptible to oxidation by atmospheric oxygen
over time. For this reason, the samples should be prepared
immediately before the titrations.
 Metaphosphoric acid is used because it serves Vitamin C in
solution against aerobic or atmospheric oxidation.
Ascorbic acid + I2 → 2I- + dehydroascorbic acid 18

19. Result
Quantity of Vitamin C in the given sample = A × F × 8.806
Here, A= Volume of 0.11 N I2 required (mL)
F= Normality factor of iodine solution=
0.10 𝑁
0.1 𝑁
= 1
Now, Quantity of Vitamin C in the given sample = 11.47 (mL) × 1× 8.806
= 101.00 g
% Error =
101.00 𝑔 −100 𝑔
100 𝑔
× 100%
= 1%
19

20. Thanks!
Do you have any questions?
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