The Equivalent Weight of the
Francesco Pantani
University of Florence
Iodate Ion in Oxidimetric Titrations
Italy
Potassium iodate is a well-known oxidising agent and is widely used in volumetric analysis. Its use in titrations is generally discussed in courses on analytical chemistry, but is not always used in the lab. The reason for this is that practical work is rather limited, and priority is given to titrations with other ions such as permanganate (e.g., Zimmermann's titration of iron), dichromate, hromate, and thiosulphate ions (e.g., the classical iodometric determination of copper). Titrations with potassium iodate can not only solve many analytical problems but are also of educational importance, since iodate ions are involved in numerous redox systems. In this connection, it is important to realize that the equivalent weight varies with the conditions, for the stoichiometry, which is the basis of volumetric methods, varies from case to case. For example, the titration of a thiosulphate solution with iodate is based on the reaction 10,-+51-+6H+-3It+3H~O
Figure 1.
or, in the case of the excess iodide, on the reaction IOa-+8I-+6Hf-31s-+
3Hs0
(2)
Any good book on analytical chemistry adds that the equivalent weight of the iodate is mole. However, the student does not always know why, for he makes i t '/, mole, saying that 1 gram-mole of iodate oxidizes 5 gram-mole of iodide. He does not normally realize that the above reaction merely makes iodine available (in a quantity of 6 g eq per mole of iodate), and the thiosulphateis in effect titratedwith iodine. The Andrews-Jamieson titrations with iodate are carried out in strong hydrochloric acid media and are based on the following reduction HIOI
+ 5 H + + 2 CI- + 4 e-
+ ICIt-
+ 3 HnO
I'o;/I-
(1)
mole ratio
Titration of I-with 101-mt various HCI consentrotionr.
the second end-point. However, this is not immaterial as regards the visual estimation of the end-point, by Andrews' method, because this relies on the formation of ICla-. We therefore have here a case where the equivalent weight changes with the method of detecting the end-point. Andrews' method involves a solvent such as CHCb, which extracts Ia as it is formed. This considerably changes the redox equilibrium, and hence the formal potential of the redox pair. This is illustrated on the redox system 12/1-in Figure 2. The diminished activity of Iain the aqueous phase, caused by the extraction,
(3)
the complete reaction, in the case of the titration of iodide, is HIO,+2I-+5H++6C1--3ICb-+3HnO
(4)
Here, the stoichiometric ratio is 1:2, and hence the equivalent weight of the iodate is of the molecular weight. It depends mainly on the acidity which of the two systems can he used. Figure 1 shows that, with hydrochloric acid more dilute than 1 N, only one endpoint is observed: the one corresponding to the formation of iodine. As the C1- concentration is increased, however, a second potential jump corresponding to IC12- begins to appear. Under certain conditions, i t is immaterial from the potentiometric point of view whether the titration is stopped after the first or after
Ill 4 N HCl, aqueous solution; (2) Figure 2. Titration of I- with Ion-: idem 1:I vol. CHCls; (31 idem 1:l vol. Cans.
+
+
Volume 47, Number 4, April 1970
/
309
reduces the formal potential and improves the potentiometric detectability of the first end-point. However, the decrease in the activity of 1% increases the formal potential of the other redox pair, namely IClZ-/ 11,whifh means that, in an acidic medium where two potentlometric jumps are observed, the improvement in the detectability of the first end-point in the presence of an extracting solvent takes place a t the expense of the detectability of the second end-point. There are also cases where the first end-point disappears altogether. This happens in acetone solution,' where iodide is directly oxidized into iodoacetone, i.e., to an oxidation state of 1
+
10,-
+ 2 I- + 3 CHaCOCHs + 3 H
+
-
3 CHICOCHJ
+ (')
The titration curve then shows only one potential jump, namelv the one corres~ondineto the second end-~oint in the previous curves. I n this case, the equivalent weight of the iodate is again '/4 of the molecular weight. Another way of suppressing the first end-point is to introduce into the hydrochloric medium a large excess of a mercuric saltZ (in practice, one may double the volume by admixing a saturated solution of HgCL). The iodocomplexes of mercury are much more stable than the chlorocomplexes, and are formed in preference to the latter. Almost all the iodide is bound in this form, and hence its activity is greatly reduced. As a result, the normal potential of the redox system, in which the reduced I- is hound, increases by a few hundred millivolts. This masks the first potential jump, hut leaves the second visible (see Fig. 3).
-
I I \
,
',
o ~i,,,.
lO-'M
50% oxidiring agent
50
, 100
4.
sped.. whish present insolution during the oxidation I-in water (left) and in a~sto"itrile(right).
wise some of the iodide is oxidized by dissolved atmospheric oxygen, thereby distorting the results. The difference between the titration in water and that in acetonitrile is due to the different stability of 13- in these two media. The stability constant K
=
111-]/[In1 [I-]
of this species is 710 in water and 2.5 X lo7in acetoniFigure 4 (left side) shows the species which are expected to he formed in an aqueous solution as the titration proceeds. The graph is based on a constant iodine concentration of g-atom/l. More than one species is seen to he present in each case; 13-, which is an oxidized and a reduced specim a t the same time, is always present. However, it never predominates greatly over I%, and SO no end-point is observed before reaching the Izstage in the oxidation. The situation is different when the titration is carried out in acetonitrile (see Fia. 4, right side). Since 13-is very stable in this medium, it is formed almost fully before the formation of Izbegins. The titration curve therefore shows an end-point for the formation of Ia- (Fig. 5).
I
112 Figure 3. Titrotion of I- with 101- in the presenceof HgCh at various HCI concentrations.
There are interesting variations in the titration curve when the solvent is varied. I n addition to water, we used 90% by volume of acetonitrile. The amount of water could not he reduced further because of the limited solubility of the salts, hut the general considerations for the present purposes are the same. It should be added that the solutions must be de-aerated, otherBERG,R., Z. anal. Chem., 69,369 (1926). N. H., AND MILLER,C. O., J. Am. Chem. Soc., 59,
"FURMAN,
152 (1937). a
DESBARRES, J., Bull. Soc. Aim. France, 502 (1961).
310 / Journal of Chemical Education
Figure 5.
N HCI,
Titration of I- wilh
103-
in 90% by volume osotonitrile
+ 0.2
The extent to which 1%is present depends of course on the iodide concentration. While Is- is virtually undetectable when the concentration in water is below M , the curves obtained with concentrations of at least 1 M resemble those ohtained in acetonitrile (Fig. 4, right side). By greatly lowering the concentrations, furthermore, these second curves may become identical with those obtained in water.
