In acid solution practically all oxidizing agents will oxidize iodide ion to iodine quantitatively. The iodine formed in the reaction can then be titrated by means of a standard sodium thiosulfate solution. This type of indirect titration is given the general name of iodometry.
Iodometric methods of analysis have a wide applicability for the following reasons:
1. Potassium iodide, KI, is readily available in high purity.
2. A good indicator, starch, is available to signal the equivalence point in the reaction between
iodine and thiosulfate. Starch turns blue-black in the presence of iodine. Therefore, when the
blue-black color disappears, the iodine has been completely reduced to the iodide ion.
3. Iodometric reactions are rapid and quantitative.
4. A precise and stable reducing agent, sodium thiosulfate
(Na2S2O3), is available to react with the iodine.
The amount of iodine liberated in the reaction between iodide ion and an oxidizing agent is a
measure of the quantity of oxidizing agent originally present in the solution. The amount of
standard sodium thiosulfate solution required to titrate the liberated iodine is then equivalent to
the amount of oxidizing agent. Iodometric methods can be used for the quantitative
determination of strong oxidizing agents such as potassium dichromate, permanganate, hydrogen
peroxide, cupric ion and oxygen.
As has been mentioned above, the endpoint in a titration of iodine with thiosulfate is signaled
the color change of the starch indicator. When starch is heated in water, various decomposition
products are formed, among which is beta-amylose which forms a deep blue-black complex with
iodine. The sensitivity of the indicator is increased by the presence of iodide ion in solution.
However, if the starch indicator solution is added in the presence of a high concentration of
iodine, the disappearance of the blue-black color is very gradual. For use in indirect methods, the
indicator is therefore added at a point when virtually all of the iodine has been reduced to iodide
ion, causing the disappearance of the color to be more rapid and sudden. The starch indicator
solution must be freshly prepared since it will decompose and its sensitivity is decreased.
However, a properly prepared solution will keep for a period of a few weeks. A preservative
such as a small amount of mercuric ions may be added to inhibit the decomposition.
Solutions of sodium thiosulfate are made up to an approximate concentration by dissolving
sodium salt in water that has previously been boiled. Boiling the water is necessary to destroy
micro-organisms which metabolize the thiosulfate ion. A small amount of
Na2CO3 is added to the
solution in order to bring the pH to about 9. The solution is standardized by taking a known
amount of oxidizing agent, treating it with excess iodide ion and then titrating the liberated iodine
with the solution to be standardized. Oxidizing agents such as potassium dichromate, bromate,
iodate or cupric ion can be employed for this procedure. You will be using potassium iodate,
KIO3, as your primary standard. The reaction between
IO3- and I- is given
Reactions Involved in Iodometric Processes
Iodometric methods depend on the following equilibrium:
Since the solubility of I2 in water is quite low, the formation of the tri-iodide ion,
I3-, allows us to
obtain useful concentrations of I2 in aqueous solutions. The equilibrium constant for
is approximately 700. For this reason iodometric methods are carried out in the presence of
excess iodide ion.
The reaction between iodine and the thiosulfate ion is:
This reaction proceeds quantitatively in neutral or slightly acidic solutions. In strongly
acidic solutions the oxidation of the thiosulfate does not proceed by a single reaction. In the
former, the thiosulfate ion is oxidized to sulfate as well as to the tetrathionate. In the latter, the
thiosulfuric acid formed undergoes an internal oxidation-reduction reaction to sulfurous acid and
sulfur. Both of these reactions lead to errors since the stoichiometry of the reactions differs from
that shown above for the thiosulfate as a reducing agent. The control of pH is clearly important.
In many cases the liberated iodine is titrated in the mildly acidic solution employed for the reaction
of a strong oxidizing agent and iodide ion. In these cases the titration of the liberated iodine must
be completed quickly in order to eliminate undue exposure to the atmosphere since an acid
medium constitutes an optimum condition for atmospheric oxidation of the excess iodide ion.
The basic reaction in the determination of copper using the iodometric method is represented
This is a rapid, quantitative reaction in slightly acidic solutions, if there is a large excess of
ion present and if the copper is in the form of a simple ion rather than a complex one. The iodine
that is liberated can be titrated in the usual manner with standard thiosulfate solution. The
reaction involving cupric ion and iodide takes place quantitatively since the cuprous ion formed as
result of the reduction is removed from the solution as a precipitate of cuprous iodide.
Iron interferes since iron(III) ions will oxidize iodide. Since the iron will be found in the +3
oxidation state as a result of the dissolution of the brass sample, a means of preventing this
interference is necessary. This can be accomplished by converting the iron(III) to a soluble
iron(III) phosphate complex using phosphoric acid. At a pH of 3.0-4.0 the iron phosphate
complex is not reduced by iodide ion. If arsenic and antimony are present, they will provide no
interference at this pH if they are in their higher oxidation states. Brass formulations also may
contain up to 39% Zn, 2.5% Sn and 8.5% Pb. When dissolved in concentrated nitric acid, the
zinc and the lead become Pb2+ and Zn2+ . These do not interfere with
the analysis of copper because they are not reduced to the Pb+ and Zn+
states by the action of iodide ion under the conditions of this experiment. The tin is oxidized to
Sn4+ by the concentrated nitric acid and after dilution and adjustment of pH this
form becomes SnO2 which is insoluble and may be observed as an inert white
precipitate at the bottom of your flask. Under these conditions the tin does not interfere with the
Sources of Error
The following are the most important sources of error in the iodometric method:
1. Loss of iodine by evaporation from the solution. This can be minimized by having a
excess of iodide in order to keep the iodine tied up as tri-iodide ion. It should also be apparent
that the titrations involving iodine must be made in cold solutions in order to minimize loss
2. Atmospheric oxidation of iodide ion in acidic solution. In acid solution, prompt titration
of the liberated iodine is necessary in order to prevent oxidation.
3. Starch solutions that are no longer fresh or improperly prepared. The indicator will then
not behave properly at the endpoint and a quantitative determination is not possible.
Click here to see images of what one can expect in this experiment.
Preparation of a 0.10 M Standard Na2S2O3
With a graduated cylinder measure out 1 liter of distilled water. Place it in your 1 liter beaker
boil the water for at least 5 minutes. Weigh out 25 g of
Na2S2O3·5H2O and 0.1 g of
Dissolve the thiosulfate in the hot water and then cool this solution with the aid of an ice bath to
room temperature. Then add the carbonate and stir until it is completely dissolved. Transfer the
solution to your plastic 1 liter bottle. When not in use store this bottle in the darkness of your
equipment cabinet as the decomposition of thiosulfate is catalyzed by light.
Potassium iodide may contain appreciable amounts of iodate ion which in acid solution will
with iodide and yield iodine. The liberated iodine would react with thiosulfate and thereby cause
the apparent molarity of the thiosulfate to be too low. The following procedure allows for the
determination of a blank correction which will properly correct for any iodate that might be
present. Prepare a solution of exactly 2.00 g of KI dissolved in 50 mL of distilled water and then
acidify the solution with 5 mL of 3 M sulfuric acid and then immediately add 5 mL of starch
indicator. If a blue-black color appears right after mixing, use the thiosulfate solution in the buret
to determine the volume of solution required to cause the color- to disappear. This volume must
be subtracted from the standardization and analyses volumes. If the potassium iodide is
completely iodate-free no color will of course develop and no blank correction is necessary.
Standardization of the Na2S2O3 Solution
Dry approximately 2 g of potassium iodate, KIO3, at a temperature of 110
oC for one hour. Weigh
to a precision of
±0.0001g three samples of the potassium iodate having weights near 0.12 g directly into
three 250 mL
Erlenmeyer flasks. Dissolve the iodate in 75 mL of distilled water. Cover the flasks with parafilm
and store them. Rinse and fill your buret with the solution. Add 2.00 g of KI to each of the
potassium iodate solutions. If a blank correction is required add
exactly 2.00 g of KI to each. If no blank
required, the exact amount of KI is not crucial but should be close to 2 g. Then add 10 mL of 1
M HCl to one of the solutions. It will turn a dark-brown color. Immediately titrate it with the
thiosulfate solution. When the color of the solution becomes very pale yellow add 5 mL of starch
indicator. Continue the titration until the blue color of the starch complex just disappears.
Follow the same procedure with each of the other two solutions, first adding the HCl then
titrating. Correct your titration data for buret error and if necessary apply the blank correction.
Calculate the molarity of the Na2S2O3 solution. Results
should agree to within 0.2% of the
average. If you do not achieve that kind of precision, titrate additional samples.
Dissolution of the Brass Sample
The following procedures in this section make use of the hot plates in the fume hoods. The solutions of dissolved brass generally have a low volume and high acid and salt concentrations. "Bumping" or little explosions of steam in the superheated liquid can occur. You don't want your hand to be close to the mouth of the flask should the solution suddenly "bump" because drops of acid (not to mention part of your sample) will fly out of the flask and possibly onto your hand. For that reason you must use your tongs to place the flasks on the hot plate and to remove them. Don't use strips of paper towel or the rubber Hot Hands because your real hand will end up being too close to the mouth of the flask.
The brass sample which you will receive does not have to be dried before use. Accurately
out three brass samples, of about 0.3 g each, directly into separate 250 mL Erlenmeyer flasks.
the fume hood add 5 mL of 6 M HNO3. Warm the solution on a hot plate in
the fume hood until
dissolution is complete. Add 10 mL of concentrated (not 3 M) H2SO4
and continue heating until
white SO3 fumes appear. It is important that you do not mistake ordinary water
vapor for SO3
fumes. It is also important at this point that the flask not be
removed from the hood. SO3 fumes
are dangerous and ought not to be inhaled. Only when the slightly denser white fumes of
observed can you be sure that all HNO3 has been removed.
NO3- will oxidize I- and hence will seriously interfere
with the procedure. Cool the flask in air for one or two minutes and then in an ice bath, then
carefully add 20 mL of distilled H2O. Boil for one or two minutes then again cool in
an ice bath. Continue to keep the flask in the ice bath and using your medicine dropper add
concentrated NH3(aq) dropwise, and with adequate mixing, until the light-blue color
solution is completely changed to the dark-blue color of the copper tetraammine complex. As
many as 400 drops (20 mL) may be required. The solution must be kept cool in an ice bath since
the reaction between the concentrated H2SO4 and concentrated
NH3 is highly exothermic. Now
add 3 M H2SO4 dropwise until the dark-blue color just disappears.
You don't have to produce a
complete disappearance of the dark blue color but you need to approach that point.
subsequent addition of phosphoric acid will lower the pH appropriately to around 3.5. If you add
too much 3M H2SO4 the pH may turn out to be sufficiently low to
cause unwanted side reactions
to occur when you reduce the Cu2+ with iodide. If you are uncertain about the
the dark blue color you may put 50 mL of 0.06 M Cu2+ in a clean 250 mL flask and
add 12 M
ammonia dropwise until you have that unmistakable dark blue color. Then add 3M
dropwise until the blue color almost disappears. Then add 2 mL concentrated
and you ought to see the dark color completely disappear. You may copy that procedure to
achieve an appropriate pH of around 3.5 for subsequent steps in the analysis. Now, back to your
real sample: Once you are confident that you haven't added too much 3M
H2SO4, but that you
have caused the dark color of the copper tetraammine complex almost to disappear, add 2.0 mL
of concentrated phosphoric acid, H3PO4, to each sample. Verify to
yourself that they exhibit the
light copper color rather than the dark color and cover the flasks with parafilm and set them aside
until you are ready to proceed with the titration.
Titration of the Dissolved Brass Sample
If you have let the dissolved samples stand overnight, be sure to warm the sample on a hot
(this can be done at your desk) in order to dissolve all larger crystals of copper sulfate that might
have formed. Be sure to cool the samples to room temperature, or below, with the aid of an ice
bath. The solutions will still contain a fine, white precipitate at this point; however, this will not
interfere with the rest of the procedure. From this point on work with only one sample at a time.
Add 4.0 g of KI to one of your samples and titrate immediately with the standard thiosulfate
solution. The sample contains white CuI precipitate and the color of
I3- must be observed against
that precipitate. The slurry will at first appear brown or dark yellow-brown. Continue adding
thiosulfate until the slurry is a light mustard color. At this point add 5 mL of starch indicator and
titrate until the mixture in the flask takes on a milky pink or lavender hue. Now add 2 g of KSCN
and mix well; the solution will darken somewhat. After the addition of thiocyanate, continue to
add more thiosulfate dropwise. You should observe a sudden change to a white or cream color.
That is the endpoint of the titration. After you have titrated all three samples calculate the
percentage of Cu in each of the brass samples, the average percentage and the average deviation.
The description above applies for brass samples with low concentrations of zinc (<10%). Some of you may have brass samples with higher concentrations of zinc. Such samples will become quite dark after the addition of KI and will lighten only slightly as thiosulfate is added. The "mustard color" will be darker than samples having low percentages of copper. When the starch is added the sample will become dark blue-black again and as you approach the end point with the thiosulfate the slurry will turn a violet color rather than milky pink or lavender hue. With the addition of the KSCN the solution will darken somewhat as in the case of the other samples, but the final end point will be a bit darker than the white or cream color described above. If you think that you have a sample with high zinc content, observe your progress carefully and take notes which will allow you to achieve repeatability.
Explanation: The reduction of Cu2+ to Cu+ occurs as the result of the oxidation of I- to I2. The I2 combines with iodide ion to produce the dark brown triodide ion, I3-. The excess iodide ion also causes the reduced copper to precipitate as white cuprous iodide, CuI. I2 and I3- in solution tend to adsorb on the surface of CuI thus becoming unavailable for rapid reduction by the thiosulfate. As a result, iodometric titrations involving reduced copper tend to yield lower results unless the adsorbed I2 can be liberated by adding thiocyanate ion, SCN-, which competes with the adsorbed iodine molecules on the surface of solid particles of CuI. After the addition of thiocyanate, continue to add more thiosulfate dropwise. You should observe a sudden change to a white or cream color. That is the endpoint of the titration. After you have titrated all three samples calculate the percentage of Cu in each of the brass samples, the average percentage and the average deviation.
Your report must include the following information in the two sections below.
1. Unknown number
2. The three weights of KIO3 used for the standardization of thiosulfate
3. Volume in mL of thiosulfate for each of your standardization titrations
4. Average molarity of the thiosulfate solution
5. Weight of brass used for each sample
6. Volume of thiosulfate solution used for each sample
7. Cu percentage for each sample
8. The average percentage of Cu in your brass sample
9. The average deviation from the mean of the percent Cu for the three samples
10. Pages in your lab notebook
containing the pertinent data
1. Why is it necessary to boil the water used to prepare the thiosulfate solution?
2. Why is Na2CO3 added to the thiosulfate solution?
3. Why is the thiosulfate solution stored in the dark?
4. Why is HCl added to the IO3- mixture and why must the solution be titrated immediately?
5. Why is the solution containing the dissolved brass sample heated to expel SO3 fumes?
6. Why is H3PO4 added to the brass sample?
7. What is the purpose of the KSCN that is added just before the endpoint in the titration?
8. Why is the solution containing the dissolved brass made basic with concentrated NH3 and then again acidified with H2SO4?
9. What is the formula of the tetrammine copper(II) complex?
10. Why do Zn2+ and Pb2+ not interfere in this procedure?
11. What sort of complications would arise if the iodinethiosulfate titration were carried out in a highly acidic solution?
12. If the solution were highly basic, how would the iodinethiosulfate reaction be influenced?
13. Why is the starch indicator not added at the beginning of the tritration?