Iodine: color, 6, 59 ions

Aside from these three classes (species with unfilled inner subshells, with unpaired electrons, or with two different oxidation states of the same element), there are a number of colored inorganic substances about which generalizations may be set up only with difficulty. Among these are many of the elementary nonmetals, a large number of covalent salts (such as mercuric iodide, cadmium sulfide, silver phosphate and lithium nitride), a number of nonmetal halides (iodine monochloride, selenium tetrachloride, antimony tri-iodide, etc.), and the colored ions, chromate, permanganate, and Ce(H20) v, whose central atoms presumably have rare-gas structures. 

Bromate and iodate salts are prepared on a much smaller scale than chlorates. Under appropriate conditions, these ions undergo oscillating chemical reactions known as chemical clocks. The best known clock reaction is observed when an acidified solution of sodium sulfite (Na2S03) is mixed with an excess of iodate in the presence of starch indicator. After a suitable induction period allowing for sodium sulfite reduction of iodate to iodide [Eq. (44)], the blue, starch-iodine color periodically appears and disappears as the iodide is oxidized to iodine [Eq. (45)], and the iodine is reduced back to iodide [Eq. (46)]. 

The behavior of antimony salts toward sulfuric acid solutions of hydroxyanthraquinones is analogous to that of boric acid they must be oxidized to antimony by means of chlorine water. Larger amounts of colored ions (Fe+ , Cu+2, Ni+, Cr+ etc.) make it difficult to detect the color change of the reagent solutions. Be+ ions lower the sensitivity of the test. Iodides interfere by releasing iodine they can be removed by precipitation with Ag2S04. Oxidizing acids impair the test. 

Bromide ndIodide. The spectrophotometric determination of trace bromide concentration is based on the bromide catalysis of iodine oxidation to iodate by permanganate in acidic solution. Iodide can also be measured spectrophotometricaHy by selective oxidation to iodine by potassium peroxymonosulfate (KHSO ). The iodine reacts with colorless leucocrystal violet to produce the highly colored leucocrystal violet dye. Greater than 200 mg/L of chloride interferes with the color development. Trace concentrations of iodide are determined by its abiUty to cataly2e ceric ion reduction by arsenous acid. The reduction reaction is stopped at a specific time by the addition of ferrous ammonium sulfate. The ferrous ion is oxidi2ed to ferric ion, which then reacts with thiocyanate to produce a deep red complex. 

B. Self-indicating reagents. This is well illustrated by potassium permanganate, one drop of which will impart a visible pink coloration to several hundred millilitres of solution, even in the presence of slightly coloured ions, such as iron(III). The colours of cerium(IV) sulphate and of iodine solutions have also been employed in the detection of end points, but the colour change is not so marked as for potassium permanganate here, however, sensitive internal... 

When iodine dissolves in organic solvents, it produces solutions having a variety of colors. These colors arise from the different interactions between the I2 molecules and the solvent (Fig. 15.21). The element is only slightly soluble in water, unless I ions are present, in which case the soluble, brown triiodide ion, I,, is formed. Iodine itself has few direct uses but dissolved in alcohol, it is familiar as a mild oxidizing antiseptic. Because it is an essential trace element for living systems but scarce in inland areas, iodides are added to table salt (sold as iodized salt ) in order to prevent an iodine deficiency. 

FIGURE 15-21 Solutions of iodine in a variety of solvents. From left to right, the first three solvents are tetrachloromethane (carbon tetrachloride), water, and potassium iodide solution, in which the brown I. ion forms. In the solution on the far right, some starch has treen added to a solution of I,. The intense blue color that results has led to the use ot starch as an indicator for the presence of iodine. 

In this step the reddish brown color of the triiodide begins to fade to yellow and finally to clear, indicating only iodide ions present. However, this is not the best procedure for determining when all of the I3 has disappeared since it is not a sensitive reaction and the change from pale yellow to colorless is not distinct. A better procedure is to add a soluble starch solution shortly prior to reaching the end point, since if it is added to soon, too much iodine or triiodide ion may be present forming a complex that may not be reversible in the titration. The amount of thiosulfate is proportional to the amount of hypochlorite ion present. 

Next, drops of bromine water, chlorine water and iodine water were mixed in samples of mineral oil and the color of the oil and evidence of a reaction were recorded in the Table 2 below. Also tested was the halide ion mixed with mineral oil. 

The hypochlorite ion may be identified most distinctly by ion chromatography. Its concentration in the aqueous solution combined as CIO and molecular CI2 (which is partly formed when hypochlorite is dissolved in water) can be measured by iodometric titration. A measured volume of sodium hypochlorite solution is added to a small volume of an acidified solution of potassium iodide (in excess). Iodine liberated is titrated with a standard solution of sodium thiosulfate or phenyl arsine oxide using starch as indicator. Blue color of starch solution decolorizes at the end point. 

The solution to the problem was discovered when a titrated sample (clear solution) was left on the bench and, after a period, it started changing back to a faint yellow color. We hypothesized that air oxidation may have caused that effect and, consequently, air may have interfered with analysis. Standard samples prepared and purposely delayed during the analysis showed that end-point volumes were larger, indicating that some of the iodide ions turned into free-iodine by air oxidation which, in turn, required more thiosulfate for titration and, therefore, larger end-point volume. The following chemical equations obtained from the literature 8) show what happens before, during, and after titration. The reaction of a chlorinated isocyanuric acid compound with potassium iodide in acidic pH is ... 

Several methods have been introduced which express the degree of oxidation deterioration in terms of hydroperoxides per unit weight of fat. The modified Stamm method (Hamm et at 1965), the most sensitive of the peroxide determinations, is based on the reaction of oxidized fat and 1,5-diphenyl-carbohydrazide to yield a red color. The Lea method (American Oil Chemists Society 1971) depends on the liberation of iodine from potassium iodide, wherein the amount of iodine liberated by the hydroperoxides is used as the measure of the extent of oxidative deterioration. The colorimetric ferric thiocyanate procedure adapted to dairy products by Loftus Hills and Thiel (1946), with modifications by various workers (Pont 1955 Stine et at 1954), involves conversion of the ferrous ion to the ferric state in the presence of ammonium thiocyanate, presumably by the hydroperoxides present, to yield the red pigment ferric thiocyanate. Newstead and Headifen (1981), who reexamined this method, recommend that the extraction of the fat from whole milk powder be carried out in complete darkness to avoid elevated peroxide values. Hamm and Hammond (1967) have shown that the results of these three methods can be interrelated by the use of the proper correction factors. However, those methods based on the direct or indirect determination of hydroperoxides which do not consider previous dismutations of these primary reaction products are not necessarily indicative of the extent of the reaction, nor do they correlate well with the degree of off-flavors in the product (Kliman et at. 1962). 

Applicable to 0.15—10 mg L1 in colored or turbid waters. Iodine and bromide interfere by forming a less soluble precipitate with the silver ions. [Fe N), CrO2-, Cr2Of, and Fe3" interfere. 

Amylose-iodine complexes have a deep blue color, which is a result of an electron relay on the polyiodide ions.196 The helix of amylose provides a tunnel for iodine molecules to align. Stability of the amylose-iodine complex has been studied.197 Iodine has been widely used for quantification of amylose contents despite the fact that the blue color development is affected by many factors, including temperature, pH and mechanical mixing. Several improved methods have been reported (see Section 6.III.1). 

Reduction of the copper(II) halides to copper(I) halides may be carried out using sulfite or copper metal. If the reduction with copper metal is carried out in concentrated HC1 solution, the solution takes on an intermediate black color before fading to the colorless CuCl ion. The black color is probably due to dimeric or polymeric ions having copper in both the +1 and +2 states, for example, Cl—Cu—Cl—CuC12(H20) . The reduction of divalent copper by cyanide to form the cyano complex, Cu(CN)73, was mentioned in Chapter 10, and the reaction between cupric and iodide ions is similar to this, releasing iodine and forming the very insoluble copper iodide ... 

Note that the product formed by reaction with pyridine (reaction c) contains the tri-iodide ion, also present in solutions containing both iodine and potassium iodide thus, solutions of iodine in pyridine and iodine in aqueous KI have about the same color. 

The basic (nucleophilic) reagents in the reactions of the preceding paragraph were the iodide and thiosulfate ions, respectively. If chlorine, bromine, or iodine is dissolved in aqueous alkali, the halogen color disappears, and an apparent self-oxidation takes place. Half of the halogen is converted to the — 1 state, the remainder to the +1 state (as a hypo-halous acid, HOX, or its ion). Such a reaction is similar to those above ... 

The following procedure is used in the analysis of iodoso and iodoxy compounds. In a 200-cc. iodine flask are placed 100 cc. of water, 10 cc. of 6 N sulfuric acid, 2 g. of iodate-free potassium iodide, 10 cc. of chloroform, and finally the sample, about 0.25 g. The flask is shaken for fifteen minutes (or longer, if the reaction is not complete), and then the mixture is titrated with 0.1 N sodium thiosulfate. If the sample is pure the change of color in the chloroform layer may be taken as the end point, but if impurities are present starch must be used, for the impurities impart a brownish color to the chloroform. This solvent is desirable, as it facilitates the reaction with potassium iodide by dissolving the reaction products. Iodosobenzene may be differentiated from iodoxybenzene, for the former reduces iodide ion in a saturated sodium borate solution, whereas the latter does not.1 The reactions involved are ... 

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