The following chart contains some of the known compounds of the halogens at various oxidation states — Physical Chemistry — Thermodynamics Chemistry Question
Some chemistry of iodine
The following chart contains some of the known compounds of the halogens at various oxidation states. Note that the oxidation number ranges in general from –I to VII, but fluorine differs significantly from the other halogens in that it has no stable oxyacids.
[VISUAL]
A glance at the chart should convince you that oxidation–reduction reactions are a very important part of halogen chemistry.
Although iodine will show some chemistry unique to itself, many of its reactions are typical of other halogens. In this experiment we shall investigate some reactions of iodine and note the influence of hydrogen ion concentration on the equilibria.
Chemicals
The following aqueous solutions are at your disposal to perform the simple experiments as given in the following Parts I, II, and III:
* potassium iodide, KI, c = 0.1 mol dm –3
* silver nitrate, AgNO3, c = 0.1 mol dm –3
* peroxide hydroxide, H2O2, 3%
* sulfuric acid, H2SO4, c = 6 mol dm –3
* potassium hydroxide, KOH, c = 6 mol dm –3
* sodium sulfite, Na2SO3, c = 0.1 mol dm –3 and c = 1 mol dm –3
* nitric acid, HNO3, c = 6 mol dm –3
Procedure
Caution: Solid iodine and its vapor will cause burns and stains on skin or clothing. Its vapors are poisonous and even small quantities will irritate the mucous membranes, if inhaled. Avoid unnecessary contact.
Use 13 × 100 mm test tubes throughout this experiment except in Part IIb.
Preliminary Experiment – The Starch Iodine Test.
Prepare a dilute solution of iodine by adding one or two small iodine crystals to about 5 cm 3 of tap water. Warm slightly, add 3 or 4 drops of starch solution, and observe. This is a very sensitive test for molecular iodine.
Note: The color is due to a starch–iodine complex which is attributed to the ability of I2 molecules to fit into the long, hollow spaces between the helical coils which constitute the starch molecule. The fit is close and the interaction strong enough to give the intense color even at very low iodine concentrations.
Part I. Some reactions of iodide ion I – .
a) Add 2 cm 3 of potassium iodide solution to 2 cm 3 of silver nitrate solution. Note the result.
b) Add 5 cm 3 of starch solution to 2 cm 3 of potassium iodide solution and finally add a drop or two of commercial bleach (5% NaOCl) solution. Note the result. Continue in adding the bleaching solution until there is a second color change. How do you account for this?
c) To 2 cm 3 of potassium iodide solution and 5 cm 3 of starch solution. Note the result.
Part II. Some reactions of iodate Ion IO3 – .
Pour about 5 cm 3 of saturated solution of KIO3 into each of two test tubes.
a) Add 3 cm 3 of KI solution and 2 cm 3 of H2SO4 solution to one of the test tubes. Decant the supernatant liquid from the solid produced. Filter, if necessary. Wash the solid with water. Do you recognize the solid? Run an identification test you have used previously to confirm your inference.
b) Add 3 cm 3 of KI solution and 2 cm 3 of KOH solution to the second test tube. What do you conclude about the role of hydrogen ion in the reaction between iodide and iodate ions?
c) Add 3 cm 3 of an acidified solution of sodium sulfite (c = 0.1 mol dm –3 ), 2 cm 3 of H2SO4 solution and 3 or 4 drops of starch solution. What do you observe?
Part III. Reaction of I2 in a basic solution.
a) To a few crystal (about 0.5 g) of solid iodine add from a dropper about 10 drops of potassium hydroxide solution. Shake the test tube gently until the solid iodine disappears and the solution is colorless. You may need to warm the solution gently and add a few more drops of KOH solution. You will identify the product of this reaction in Part d.
b) Cool the solution and make it acidic by adding sufficient (10 drops or slightly more) of the HNO3 solution to neutralize the base added previously. Note the product of this reaction. What do you think it is?
c) Make the solution basic again by adding a few drops of KOH solution. Warm gently and add a few drops of KOH solution more, if necessary, until a color change is observed. Discard the solution.
d) Repeat the procedure outlined in Part a. Cool under the cold water tap until a solid crystallizes from the solution. Decant the supernatant liquid and save it for part (2) below.
1. Dry the white solid by heating the test tube gently. Allow it to cool. Dissolve the white solid in 5 cm 3 of water. Add 5 cm 3 of sodium sulfite solution (c = 1 mol dm –3 ) 2 cm 3 of H2SO4 solution and 3 or 4 drops of starch solution. Note the result. Compare it with that obtained in Part IIb.
2. To the decanted liquid add 5 – 10 drops of AgNO3 solution; shake the test tube and note the result. Compare the product with that obtained in Part Ia.
Write the equations for the reactions observed in Parts Ia, Ib, Ic.
Model Answer
I a: I – (aq) + Ag + (aq) AgI(s) (yellow precipitate)
I b: 2 I – (aq) + OCl – (aq) + 2 H + (aq) I2 + Cl – (aq) + H2O(l)
The solution was deep blue colored . The color is due to the starch-iodine complex.
I2 + OCl – (aq) + H2O(l) IO3– (aq) + 5 Cl – (aq) + H + (aq)
In excess of NaOCl the iodine was further oxidized to iodate.
I c: 2 I – (aq) + H2O2(aq) + 2 H + (aq) I2 + 2 H2O(l)
The solution was deep blue colored. The color is due to the starch-iodine complex.
a) How did the results in Part IIId(1) compare with those obtained in Part IIb?
b) How did the test with 0.1 M silver nitrate in Part IIId(2) compare with the results of Part Ia?
c) What do you conclude about the ionic species formed when I2 reacts with KOH (6 mol dm –3 ) as in Part IIIa?
Model Answer
a. In Part IIb the sulfite ions (SO3 2– ) was reacted with an excess of iodic ions (IO3 – ) (saturated solution of KIO3) and in the presence of starch indicator the deep-blue starch-iodine color increased systematically as a result of the following of reactions:
IO3 – + 3 SO3 2– I – + 3 SO4 2–
5 I – + IO3 – + 6 H + 3 I2 + 3 H2O
In Part IIId the iodate ions (IO3 – ) are reacted with an excess of sulfite ions (SO3 2– ). With the excess of sulfite, free iodine periodically appears and disappears as a result of the following sequence of reactions:
IO3 – + 3 SO3 2– I – + 3 SO4 2–
5 I – + IO3 – + 6 H + 3 I2 + 3 H2O
3 I2 + 3 SO3 2– + 3 H2O 6 I – + 6 H + + 3 SO4 2–
The net reaction is the oxidation of iodates to iodides and the starch indicator oscillates between deep blue and almost colorless as the iodine concentration pulsates.
b. In Part IIId(2) the product is the same (yellow precipitate) as that obtained in Part Ia, following the reaction:
I – (aq) + Ag + (aq) AgI(s)
c. The anionic species formed when I2 reacted with 6 M KOH as in Part IIIa was the iodates (IO3 – ) and iodides (I – ) anions.
Write the equation for the self-oxidation-reduction reaction of iodine in a basic solution. Write the equation for the reverse reaction in an acid solution.
Model Answer
The equation for the self–oxidation–reduction reaction of iodine in a basic solution is:
3 I2 + 6 OH – 5 I – + IO3 – + H2O
and the reverse of this reaction in an acid solution:
5 I – + IO3 – + 6 H + 3 I2 + 3 H2O
In which oxidation state do the halogens occur most commonly in nature? Explain your answer in terms of the electronic structure of this species for chlorine.
Model Answer
As X –, oxidation state (–I), as the result of their valence electronic configuration:
…..ns 2 np 5 .
How would you prepare elemental fluorine, F2? Consult an oxidation-reduction table to check the feasibility of your method.
Model Answer
The only practicable method of preparing F2 gas is based on the electrolysis of fluoride salts, i.e., potassium fluoride (KF) dissolved in anhydrous HF:
KF + HF F2 + H2 (electrolysis)
2 F – F2 + 2 e – Eo = – 2.87 V
2 H + + 2 e – H2 E o = 0.00 V
2 F – + 2 H + F2 + H2 E = – 2.87 V
Find the geometry, using the VSEPR model, for the following anions of the halogen oxoacids: ClO2 –, ClO4 –, BrO3 –, IO6 5–.
Model Answer
ClO2 – : bent
ClO4 – : tetrahedral
BrO3 – : trigonal pyramidal
IO6 5– : octahedral