CHAPTER I

INRODUCTION

A. Objective
Students are able to:
1. Prepare primary and secondary standard solution for iodometry titration
2. Perform iodometry titration and observe the changes that occur at the end of the
titration.
3. Calculate the percentage of vitamin C in a drinking sample
4. Use pycnometer to determine solution density

B. Literature Review
B.1. Iodometry

The term “iodometry” describes the type of titration that uses a standardised sodium
thiosulfate solution as the titrant, one of the few stable reducing agents where oxidisation of air
is concerned. Iodometry is used to determine the concentration of oxidising agents through an
indirect process involving iodine as the intermediary. In the presence of iodine, the thiosulphate
ions oxidise quantitatively to the tetrathionate ions. (5) Iodometry reaction has low, pH
independent redox potential, and reversibility of the iodine/iodide reaction, iodometry can be
used both to determine amount of reducing agents (by direct titration with iodine) and of
oxidizing agents (by titration of iodine with thiosulfate). The method of end point detection is
based on blue starch complex.(3)

Reversible iodine/iodide reaction:

2I- ↔ I2 + 2e-
The determination whether it should be treated as oxidation with iodine or reduction with
iodides depends on the other redox system involved.(3)
In case of iodometry, reducing agent (sodium thiosulphate) is titrated directly with
standard I2 solution which is produced during a reaction (i.e. reaction of some oxidizing agents
like KMnO4, H2O2 with KI). Below is the reaction
1
 2S2O32- + I2 → S4O62- + 2I-
In the case of both reactions it is better to avoid low pH. Thiosulfate is unstable in the
presence of acids, and iodides in low pH can be oxidized by air oxygen to iodine. Both processes
can be source of titration errors. (3)
The use of iodine as a titrant suffers from two major disadvantages. First, iodine is not
particularly soluble in water, and second, iodine is somewhat volatile. Consequently, there is an
escape of significant amounts of dissolved iodine from the solution. Both of these disadvantages
are overcome by adding iodide (I–) excess to iodine (I2) solutions. In the presence of iodide,
iodine reacts to form triiodide (I3–) which is highly soluble and not volatile. (5)
I2 + I– = I3–
The major chemical species present in these solutions is triiodide. The reduction of
triiodide to iodide is analogous to the reduction of iodine. (6)
I3– + 2 e– = 3 I–
Triiodide reacts with thiosulfate to yield iodide and tetrathionate. (6)
2 S2O32– + I3– = S4O62– + 3 I–
This lowers free iodine concentration and such solutions are stable enough to be used in
lab practice. Still, it should remember that their shelf life is relatively short (they should be kept
tightly closed in dark brown bottles, and standardized every few weeks). Iodine solutions are
prepared dissolving elemental iodine directly in the iodides solution. Elemental iodine can be
prepared very pure through sublimation, but because of its high volatility it is difficult to weight.
Thus use of iodine as a standard substance, although possible, is not easy nor recommended.
Iodine solutions can be easily normalized against arsenic (III) oxide (As2O3) or sodium
thiosulfate solution. It is also possible to prepare iodine solutions mixing potassium iodide with
potassium iodate in the presence of strong acid:(3)
5I- + IO3- + 6H+ → 3I2 + 3H2O
Potassium iodate is a primary substance, so solution prepared this way can have exactly
known concentration. However, this approach is not cost effective and in lab practice it is much
better to use iodate as a primary substance to standardize thiosulfate, and then standardize iodine
solution against thiosulfate. (7)
Iodine in water solutions is usually colored strong enough so that its presence can be
detected visually. However, close to the end point, when the iodine concentration is very low,

2
its yellowish color is very pale and can be easily overlooked. Thus for the end point detection
starch solutions are used. Iodine gets adsorbed on the starch molecule surface and product of
adsorption has strong, blue color. Exact mechanism behind adsorption and color change is not
known, see for example this explanation of starch as an indicator usage.(7)
In the presence of small amounts of iodine adsorption and desorption are fast and
reversible. However, when the concentration of iodine is high, it gets bonded with starch
relatively strong, and desorption becomes slow, which makes detection of the end point
relatively difficult. However, high concentrations of iodine are easily visible, so if thiosulfate is
used to titrate solution that initially contains high iodine concentration, the titration can be
carried out until the solution gets pale and add starch close to the end point. This will result in
color turning from dark blue to a colorless solution. This means that end point has been already
achieved with both reactants completely in equilibrium. (6)

B.2. Iodimetry

The term iodimetry, on the other hand, refers to titration using an iodine solution and is
useful for determining substances that have reducing properties. The half-reaction is as
follows: (3)

I3- + 2e- ↔ 3I- E0 = 0.536 V

Standard iodine solutions are of fairly limited use compared to oxidants because of
their small electrode potential. This characteristic of the I3- /I- pair can sometimes be an
advantage, however, because it makes it selective and therefore means that strong reducing
agents can be determined in the presence of weak ones. If a standard iodine solution is used as
a titrant for an oxidizable analyte, the technique is iodimetry. (3)

In iodimetry titration, iodine is used as an oxidizing agent, but only a substance that is strong
enough as an element of reduction that is directly titratedwith iodine. Therefore, the number of
iodimetric determinations is a little. Important substances that are strong enough as elements of

3
reduction to be titrated directly with iodine i.e. substances with a reduction potential
much lower is thiosulfate, arsenic (III), antimony (III), sulfide, sulfite, tin
(II) and ferocyanide, these substances react completely and quickly with iodine even
in acidic solution. With a rather weak reducing agent, for example trivalent arsenic
or trivalent stibium, a complete reaction will only occur if the solution is maintained
remain neutral or very slightly acidic, under these conditions the reduction potential of the
substance reducing agent is the minimum or the reduction power is maximum.(1)
The iodine, as oxidizing agent, needs to be maintained in terms of its pH. The pH should be
maintained at 8. The pH is maintained by the addition of NaHCO3. The bubbling action of CO2 would
remove the dissolved O2 from the solution. In addition, the CO2 forms a blanket over the solution which
prevents the air oxidation of the I-. (4)
In iodimetry (where I2 or to be precise I3- is added drop by drop to a solution containing
the reducing analyte), starch can be added at the very beginning of the titration. The first excess
drop of I3- after the equivalence point, causes the solution to turn dark blue. The end point can
bedetermined wehen there is a color change from dark blue to a certain color depending on the
sample .(4)

B.3. Indicator
Starch is used as an indicator for iodine. In a solution with noother colored species, it is possible
to see the color of 5 M I3-. With starch, the limit ofdetection is extended by about a factor of
10.In iodimetry (titration with I3-), starch can be added at the beginning of the titration.The first
drop of excess I3- after the equivalence point causes the solution to turn dark blue (4).
In iodometry (titration of I3-), I3= is present throughout the reaction up to the equivalence
point. Starch should not be added until immediately before the equivalence point (as detected
visually, by fading of the I3-). Otherwise some iodine tends to remain bound tostarch particles
after the equivalence point is reached.Starch-iodine complexation is temperature dependent. At
50°C, the color is only one-tenth as intense as at 25°C. If maximum sensitivity is required,
cooling in ice water is recommended2 Organic solvents decrease the affinity of iodine for
starch and markedly reduce the utility of the indicator.

4
 CHAPTER II
EXPERIMENTAL METHOD

A. Materials and Apparatus

A.I. Materials
1. KIO3 (Potassium Iodate)
2. Na2S2O3 (Sodium Thiosulphate)
3. Soluble starch
4. Concentrated H2SO4
5. KI (Potassium Iodide)
6. I2 (Iodine)
7. Drinking sample containing vitamin C

A.2. Apparatus
1. Weighing bottle
2. Beaker glass
3. Mixing rod
4. Funnel
5. Volumetric flask
6. Measuring cup
7. Watch glass
8. Volumetric pipette
9. Iodine Flask
10. Burette
11. Burette stand and clamp
12. Spray bottle
13. Pycnometer
14. Thermometer

5
B. Solution Preparation
B.1.1. Potassium Iodate (KIO3) ± 0,01 N Solution
Volume = 0.1 L
Molecular Weight (MW) = 214 gr/ mol
Valence =6
The Formula Equation Tolerance ± 10%
N = M x Valence 0,0356 x 10% = 0,00356
𝑀𝑜𝑙
N= 𝑥 𝑉𝑎𝑙𝑒𝑛𝑐𝑒 0,0356 + 0,00356 = 0,0392 gr
𝑉𝑜𝑙𝑢𝑚𝑒(𝐿)
𝑚𝑎𝑠𝑠 (𝑔𝑟) 0,0356 - 0,00356 = 0,0321 gr
N = 𝑔𝑟 𝑥 𝑉𝑎𝑙𝑒𝑛𝑐𝑒
𝑀𝑊 ( ) 𝑥 𝑣𝑜𝑙𝑢𝑚𝑒 (𝐿)
𝑚𝑜𝑙

𝑚𝑎𝑠𝑠 (𝑔𝑟)
0,01 N = 𝑔𝑟 𝑥6
214 𝑥 0,1 𝐿
𝑚𝑜𝑙

Mass = 0,0356 gr

B.1.2. Preparation of Potassium Iodate (KIO3) ± 0,01 N Solution

1. KIO3 powder was weighed within the range of (0,0321-0,0392) grams using analytic
balance.
2. KIO3 powder was dissolved in distilled water in beaker glass below 100 ml, the solution
was stirred using mixing rod until dissolved completely.
3. The KIO3 solution was poured into volumetric flask (100 ml) using funnel that was given
a filter paper, distilled water was added until it reached 100 ml scratch.
4. The volumetric flask was shaken until completely homogenous.

B.2.1. Sodium Thiosulphate (Na2S2O3.5H2O) ± 0,01 N Solution

Volume =1L
Molecular Weight (MW) = 248.18 gr/ mol
Valence =1
The Formula Equation
N = M x Valence
𝑀𝑜𝑙
N = 𝑉𝑜𝑙𝑢𝑚𝑒(𝐿) 𝑥 𝑉𝑎𝑙𝑒𝑛𝑐𝑒

6
 𝑚𝑎𝑠𝑠 (𝑔𝑟)
N = 𝑔𝑟 𝑥 𝑉𝑎𝑙𝑒𝑛𝑐𝑒
𝑀𝑊 ( ) 𝑥 𝑣𝑜𝑙𝑢𝑚𝑒 (𝐿)
𝑚𝑜𝑙

𝑚𝑎𝑠𝑠 (𝑔𝑟)
0,01 N = 𝑔𝑟 𝑥1
248.18 𝑥1𝐿
𝑚𝑜𝑙

Mass = 2,481 gr ≈ 2,5 gram

B.2.2. Preparation of Sodium Thiosulphate (Na2S2O3.5H2O)

1. Na2S2O3 powder was weighed around 2,5 grams with hard balance
2. Na2S2O3 powder was dissolved in 1L distilled water in beaker glass
3. Na2S2O3 solution was poured into burette until 50 mL

B.3.1. Starch Solution 1%, 150 mL

The Formula Equation
𝑚𝑎𝑠𝑠 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑒
%𝑚𝑎𝑠𝑠 = × 100%
𝑚𝑎𝑠𝑠 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛
𝑚𝑎𝑠𝑠
%𝑚𝑎𝑠𝑠 = × 100%
𝑉×𝜌
%𝑚 × 𝑉 × 𝜌
𝑚𝑎𝑠𝑠 =
100%
1% × 150 × 1
𝑚𝑎𝑠𝑠 =
100%
𝑚𝑎𝑠𝑠 = 1.5 𝑔ram

B.3.2. Preparation of Starch Solution

1. Starch powder was weighed around 1,5 grams with hard balance
2. Starch powder was dissolved in beaker glass until 150 mL
3. The solution was stirred while it was heated

B.2.1. Sulphuric Acid solution ± 2 N (H2SO4)

MW = 98 g/mol
Volume = 0.03 L
Valence = 2
Density = 1.84 g/ml

7
% purity = 96%
The Formula Equation
N = M x Valence
𝑀𝑜𝑙
N = 𝑉𝑜𝑙𝑢𝑚𝑒(𝐿) 𝑥 𝑉𝑎𝑙𝑒𝑛𝑐𝑒
𝑚𝑎𝑠𝑠 (𝑔𝑟)
N = 𝑔𝑟 𝑥 𝑉𝑎𝑙𝑒𝑛𝑐𝑒
𝑀𝑊 ( ) 𝑥 𝑣𝑜𝑙𝑢𝑚𝑒 (𝐿)
𝑚𝑜𝑙

𝑚𝑎𝑠𝑠 (𝑔𝑟)
2 N = 98,08𝑥 0,03 𝐿 𝑥 2

Mass = 2,49 gram

𝑚𝑎𝑠𝑠 × 100%
𝑉=
𝜌 × %𝑝𝑢𝑟𝑖𝑡𝑦
2.94 × 100%
𝑉=
1.84 × 96%
𝑉 = 1.664 𝑚𝐿

B.4.2. Preparation of Sulphuric Acid solution ± 2 N (H2SO4)

1. 1,644 mL of concentrated H2SO4 was taken using measuring cylinder
2. The solution was dissolved in beaker glass until 50 mL with aquadest

B.5.1. Potassium Iodide 10% solution

Volume = 30 ml
Density = 1 g/ml
%purity = 10%
Formula equation
%𝑝𝑢𝑟𝑖𝑡𝑦 × 𝑉 × 𝜌
𝑚𝑎𝑠𝑠 =
100%
10% × 30 × 1
𝑚𝑎𝑠𝑠 =
100%
𝑚𝑎𝑠𝑠 = 3 𝑔ram

B.5.2. Preparation of Potassium Iodide 10% solution

1. KI powder was weighed around 3 grams with hard balance

8
 2. The solution was dissolved in beaker glass until 30 mL with aquadest

B.6.1. Iodine (I2) ± 2 N Solution

MW = 253.81 g/mol
Volume = 0.5 L
Valence = 2
The Formula Equation
N = M x Valence
𝑀𝑜𝑙
N = 𝑉𝑜𝑙𝑢𝑚𝑒(𝐿) 𝑥 𝑉𝑎𝑙𝑒𝑛𝑐𝑒
𝑚𝑎𝑠𝑠 (𝑔𝑟)
N = 𝑔𝑟 𝑥 𝑉𝑎𝑙𝑒𝑛𝑐𝑒
𝑀𝑊 ( ) 𝑥 𝑣𝑜𝑙𝑢𝑚𝑒 (𝐿)
𝑚𝑜𝑙

𝑚𝑎𝑠𝑠 (𝑔𝑟)
2 N = 253.81𝑥 0,5 𝐿 𝑥 2

Mass = 0,6345 gram

B.6.2. Preparation of Iodine (I2) ± 2 N Solution

1. 0,6345 grams of I2 powder was weighed and 2 grams of KI was weighed, both using
hard balance
2. Both powder was dissolved in beaker glass until 0,5 L with aquadest

C. Experimental Procedure
C.1. Standardization of Sodium Thiosulphate (Na2S2O3.5H2O) Solution with Potassium
Iodate (KIO3) solution
1. 10 mL of KIO3 solution was taken using volumetric pipette then moved into iodine flask
2. 2 mL of diluted H2SO4 was added and 8 mL of KI 10% solution was also added
3. Na2S2O3 solution was titrated until KIO3 solution became light yellow
4. 3 mL of starch solution was added
5. Na2S2O3 solution was titrated again until dark blue colored solution was achieved
6. The volume of Na2S2O3 needed and the color change were recorded
7. The procedure was repeated once again

9
C.2. Standardization of Iodine (I2) ± 2 N Solution with Sodium Thiosulphate (Na2S2O3)
solution
1. 10 mL of I2 solution was taken using volumetric pipette then moved into iodine flask
2. The solution was titrated with Na2S2O3 solution until it became light yellow
3. 3 mL of starch solution was added
4. Na2S2O3 solution was titrated again until dark blue colored solution was achieved
5. The volume of Na2S2O3 needed and the color change were recorded
6. The procedure was repeated once again

C.3. Determination of Vitamin C percentage in drinking sample

1. 10 mL of drinking sample was taken using volumetric pipette then moved into iodine
flask
2. 5 mL of starch solution was added and titrated with I2 solution until dark blue color
vanished
3. The volume of I2 needed and the color change were recorded
4. The procedure was repeated once again

C.4. Determination of Density

1. The pycnometer and its lid were cleaned first, and weighed using analytic balance (m1)
2. The pycnometer was filled with aquadest until it was full and contained no bubble
3. Pycnometer lid was closed and the outside wall was dried from the remaining liquid
4. The pycnometer was weighed using analytic balance (m2)
5. The pycnometer was emptied then filled with drinking sample
6. The pycnometer was weighed with analytic balance (m3)

10
 CHAPTER III
RESULT AND DISCUSSION
A. Results
Experimental Result
A.1. Standardization of Na2S2O3 Solution with Standard Solution of KIO3
1. Standard solution was prepared by weighing 0,0351 gr of potassium iodate, then
dissolved and diluted with distilled water until 100 mL
Chemical formula of potassium iodate: KIO3
Molecular weight of potassium iodate: 214 gr/mol
2. Titration Result
Indicator: starch

Volume of
Volume of KIO3
Na2S2O3 Color Change
solution(mL)
solution(mL)
10 9,2 Red-Yellow-Black-
Colorless
10 9,0 Red-Yellow-Black-
Colorless
Average = 10 9,1 Red-Yellow-Black-
Colorless

A.2. Standardization of I2 solution using Na2S2O3 solution

1. I2 solution was prepared by weighing 0,6 gr iodine, added with 0,2 gr potassium iodide,
then dissolved and diluted with distilled water until 500 mL.
2. Titration Result
Indicator: starch

Volume of I2 solution Volume of Na2S2O3

Color Change
(mL) solution(mL)
10 6,2 Red-Yellow-Black-
Colorless

11
 10 6,4 Red-Yellow-Black-
Colorless
Average = 10 6,3 Red-Yellow-Black-
Colorless

A.3. Determination of vitamin C percentage in liquid sample

Name of sample : Oonami C
Vitamin C percentage according to the nutrition fact : AKG 160%
Indicator: Starch

Volume of drinking Volume of I2

Color Change
sample(mL) solution(mL)
10 7,7 Yellow-colorless-dark
cyan
10 7,4 Yellow-colorless-dark
cyan
Average = 10 7,55 Yellow-colorless-dark
cyan

A.4. Determination of vitamin C percentage in solid sample

Name of sample : Redoxon
Vitamin C percentage according to the nutrition fact: AKG 1111%
Indicator: starch

Volume of solid Volume of I2

Color Change
sample(mL) solution(mL)
10 42,1 Yellow-colorless-dark
cyan
10 41,4 Yellow-colorless-dark
cyan
Average = 10 41,75 Yellow-colorless-dark
cyan

12
A.5. Density determination
Aquadest temperature = 28 ⁰C
Volume of pycnometer = 10 mL
Mass of pycnometer(m1) = 16,6980 gr
Mass of pycnometer + aquadest (m2) = 27,6415 gr
Mass of pycnometer +sample(m3) =27,4514 gr

B. Calculation
B.1. Standardization of Na2S2O3 with KIO3
Equivalent of Na2S2O3 = Equivalent of KIO3
(N x V) Na2S2O3 = (N x V) KIO3
N x 9,1 = 9,8 x 10-3 x 10
NNa2S2O3 = 0,0108 N
B.2. Standarization of I2 with Na2S2O3
Equivalent of I2 = Equivalent of Na2S2O3
(N x V) I2 = (N x V) Na2S2O3
(0,0108 𝑥 6,3)
NI2 = 10

NI2 = 6,804 x 10-3 N

B.3. Determination of vitamin C mass in liquid sample
Equivalent of I2 = Equivalent of Vitamin C
(N x V) I2 = (N x V) Vitamin C
6,804 x 10-3 x 7,55 = N x 10
NVit C = 5,13702 x 10-3 N

𝑁 𝑥 𝑀𝑊 𝑥 𝑉𝑜𝑙𝑢𝑚𝑒
Mass of Vitamic C = 10
5,13702 𝑥 10−3 𝑥 176 𝑥 0,1
= 2

= 0,0452 gr (in 100 mL)

120
Mass of Vitamin C in 1 bottle = x 0,0452
25

= 0,216 gr = 216 mg

13
 216−145
Deviation = 𝑥 100%
145

= 48,9 %
B.4. Determination of vitamin C mass in solid sample
Equivalent of I2 = Equivalent of Vitamin C
(N x V) I2 = (N x V) Vitamin C
6,804 x 10-3 x 41,75 = N x 10
NVit C = 0,0284 N
MVit C = 0,01420335 M
𝑚𝑎𝑠𝑠(𝑔𝑟) 1000
0,01420335 = 𝑥
176 100

Mass of Vitamic C = 0,2499 gr

= 249,9 mg
4,5255
Mass of Vitamin C in 1 tablet = x 249,9
1,160

= 974,9 mg
974,9−1000
Deviation = 𝑥 100%
1000

= 2,51 %

C. Discussion
The experiment is done using iodine flask not Erlenmeyer because, acid solutions of iodide
are oxidised by oxygen from the air(2).

In neutral solution , the oxidation of iodide can be neglected and the rate of oxidation
increases rapidly with decreasing pH. The reaction that happened, can be catalysed with a certain
metal ions and a strong light. Therefore, the titrations should not be performed in direct sunlight
and the solution containing iodide should be stored in dark glasses(2).

The first titration is the standardization of Na2S2O3 and KIO3. The first reaction that
happened is between KIO3 and of KI in the iodine flask. In the reaction of KIO3 and KI, sulfuric

14
acid was added to the reaction. Sulfuric acid is needed, because iodate only react in acidic
environment, it will not react in neutral or low acidity environment. The colour will be reddish-
brown. Then it is titrated with Na2S2O3. Iodine actually can act as its own indicator. A solution
of iodine in aqueous iodide has an intense yellow to brown colour. The end point is made easier
to be detected by the use of a solution of starch as indicator. Starch reacts with iodine in the
presence of iodide to make the end point detection much easier. The starch solution is added
when the titration is near the equivalence point when the colour begins to fade. This is because
at high concentration of iodine it gives a water-insoluble complex with the iodine. The color
sensitivity is also affected by temperature, at 50⁰C the sensitivity is ten times less sensitive than
at 25 ⁰C. After the addition of starch the solution will become dark-blue. The titration is
continued until the solution become colorless meaning that there is no iodine that is present(2).
The second titration is the standardization of I2. The Na2S2O3 that has been
standardized in the first titration is used as the primary solution to standardize I2. In the
preparation of I2 solution, as I2 is not soluble in water. It needed the addition of excess KI in
order to be dissolved. At first the color will be reddish brown. After some addition of Na2S2O3
when the color become light yellow, the starch is added as the indicator. The color will be
dark-blue and the titration will continue until it become colorless which signal the end point(1).
In order to determine the vitamin C in liquid and solid sample, a few part of the sample
is taken and is used to be stardarized with I2 which have been standarized in the previous
experiment. Unlike the first and second titration, the starch indicator is added before the titration
start. It is because I2 is used as the titrant, and an excess of I2 in the sample will react with the
indicator causing the solution to be dark-blue which signals the end point. After knowing the
mass of the vitamin C in the few part of the sample. It is then used to determine the vitamin C
in one bottle for the liquid and 1 tablet for the solid(2).

15
 CHAPTER IV
CONCLUSION AND SUGGESTION
A. Conclusion

The mass of vitamin C of the solid sample in 1 tablet and liquid sample in one bottle
obtained from the experiment are 974,9 mg and 216 mg, while the nutrition fact says
145 mg of vitamin C in the liquid sample and 1000 mg in the solid sample.

B. Suggestion
1. Remember to be precise at expelling the titrant to the sample. Any indirect contact
between two fluid will reduce the accuracy of the result.
2. Providing more solution needed for the titration, so that the experiment can go well

16
 REFERENCES

1. Vogel’s, 1990, “ A textbook of Quantitative Chemical Analysis”. 6th ed, Jakarta :

Erlangga
2. Iodometri-Iodimetri. (n.d.). Chen, C.-Y., Yang, X.-Q., Li, J.-Q., & Zhang, M. (2017).
Effect of Na 2 S 2 O 3 and Na 2 S on the Corrosion Behavior of Q235 Steel in Sodium
Aluminate Solution , 4, 551–557. https://doi.org/10.1142/9789813226517_0080
3. iodometric-titration @ www.titrations.info. (n.d.). Retrieved from
http://www.titrations.info/iodometric-titration
4. Harris, D. C. (2015). Quantitative chemical analysis, 9th ed. Analytical and
Bioanalytical Chemistry (Vol. 407). https://doi.org/10.1007/s00216-015-9059-6
5. b85d0f0970ff93ddb02af1003fb5bbfb0fc9b69d @ www.federica.unina.it. (n.d.).
Retrieved from http://www.federica.unina.it/agraria/analytical-chemistry/iodometry/
6. Brown, S. D., & Beebe, T. P. (2005). Preparation of Standard Sodium Thiosulfate
Solution and Determination of Hypochlorite in a Commerical Bleach Product.
University of Delaware, 4. Retrieved from
https://www1.udel.edu/chem/beebe/Chem120/Chem120LAB
3HypochloriteinBleach.pdf
7. iodometric-titration-end-point-detection @ www.titrations.info. (n.d.). Retrieved from
http://www.titrations.info/iodometric-titration-end-point-detection

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