THE OXIDATION OF NITRITE AND IODATE IONS BY HYPOCHLORITE IONS1

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1 THE OXIDATION OF NITRITE AND IODATE IONS BY HYPOCHLORITE IONS1 ABSTRACT The oxidatio~~ of nitrite ions and of iodate ions by hypochlorite ions in aqueous solutio~~ has bee11 examined. The osidation of nitrite is really a reaction of hypochlorous acid, with the slow stage HOCl + NO.- + H?O + H3Of + CI- + NOa-. The rate constant is given by log k = /RT (time in minutes, and the activation energy in calories). The oxida- tion of iodate is chiefly a reaction of hypochlorite ions, probably C CI , although the rate is somen~hat increased by a higher concentration of hydroxide ions. The rate constant isgiven by log k = lg.l5-2g,loo/rt. These results are compared with other oxidations by hypochlorite ions, to see if any general trends are apparent. Sodium hypochlorite solution will oxidize many substances, and the mechanisms of these oxidations have been the subject of considerable investigation, e.g. ref. 14. Some of these reactions have, in fact, turned out to be reactions of free hypochlorous acid, but at present it does not seen1 possible to predict whether ally particular reaction will go through hypochlorite ions or the free acid. On the other hand, there does seem to be a rough correlation between the heats of these reactions and their rates; for instance, the very exothernlic oxidation of sulphite ions is very fast, while the slightly endothermic oxidation of iodate ions is slow. This is certainly what might be expected; however, there are one or two notable exceptions, and this matter will be briefly discussed later. The present work was undertaken in order to increase the i~lforn~ation on which any gener a 1' lzations about the reactions of hypochlorites could be based. 1. REACTION OF NITRITE AND HYPOCHLORITE IONS This reaction has been lcnown in relatively- acid solution for a long time, and is the basis of one method for esti~nating nitrites. Some ~neasure~nents on the reaction under these conditions have been made by Shilov (j), who found a considerable dependence of the rate on ph over the range.5 to 8. However, he gives very little detail about his results. The present worlc was on the reaction in alkaline solution, where it was thought that a reaction of two ions might occur. AS will be seen in what follows, this also proved to be a reaction of hypochlorous acid. Experimental Method Mixtures in solution of sodiunl hypochlorite, sodium nitrite, and sodium hydroxide of various concentrations were made up; these were then kept at constant temperature, and analyzed at intervals as described below. The sodiu~n hypochlorite was made in the usual way from sodium hydroxide and chlorine. The solutions consequently contained sodium chloride in amounts at least equimolar to the sodium hypochlorite. Sodiu~n nitrite and hydroxide stock solutio~ls were made from reagent grades of these chemicals. Reagent grade sodium chloride was added to control the ionic strength. The solutio~ls were kept stirred in a water thermostat, and, at intervals, sanlples were pipetted out and analyzed as follows. The cooled sample was diluted to a known '~Manzrscript received April IS, Contribzrtion frowz tlte Depa~tntent of Clte~ltistry, University of Toronto, Toronto, Ontario. Can. J. Chem. Vol. 39 (19G1) 164.5

2 1646 CANADIAN JOURNAL OF CHEMISTRY. VOL volunle in a volumetric flask, and the optical density of the diluted solution was measured at 292 mp, where there is a ~naximum in the absorption spectru~n of hypochlorite ions, and at 360 mp, which is close to the maximum in the absorption spectrum of nitrite ions. A Beckmann DU spectrophotometer was used. The known extinction coefficients for the various ions at these wave lengths were checked on separate samples, and the following values were used in the calculations: OCI NOz Nos It was assumed that all nitrogen was present either as nitrite or nitrate; hence, if the original concentration of nitrite is l<nown, measurements at these two wave lengths suffice to determine the concentrations of hypochlorite, nitrite, and nitrate ions. This assumption was justified by the consistency of the results. Most of the disappearance of hypochlorite ions is due to reaction with nitrite, but a small part (between 0.1 and 0.5%) is due to their decomposition to chlorate and chloride ions. This is shown by the slow increase of the difference, [NO2-]-[OCI-]. Results from a typical run are given in Table I. The TABLE I (Run 1; at 60 C; ionic strength 1.46; [NaOH] = M) Time, [OCl-], [SO?-], Difference, ln~n M -11 At log [OCI-]/[NO?-] last column gives log [OCI-]/[Not-], which should give a straight line when plotted against tiine if the reaction is first order in hypochlorite and in nitrite ions. This was found to be the case for all runs. The evaluation of the rate constants was clone as follows. Two reactions are occurring simultaneously in the mixture: where the rate constants are respectively kl and k?. The second reaction is followed by further oxidation of chlorite to chlorate ions. Then if [OCl-] = x, and [NO*-] = y, these equatioils require that

3 LISTER AND ROSEKBLUM: OSIDATIOX OF SITRITI*: AND IODATE IONS 1647 The rigorous solution of these equations is difficult; but since hz is small, an approximate solution can be obtained as follows. If k2 = 0, it is easily shown that the solution is where xo and yo are the initial concentratio~ls. Elder these conditions a plot of In(x/y) against time would give a straight line of slope kl(.2-.0-yo). If kz is not zero, the equations above give and since (x-y) is very nearly constant and k? is small, in practice a plot of In(x/y) against time does in fact give a virtually straight line. Table I, for instance, sho\\rs that in run 1, (x-y) varied by only 0.13%; and the slope of the graph is such that k, is of the order of 100 times as large as kr. If s is the slope of the graph of ln(.v/y) against time, then approximately where the bar over the symbols indicates the average value during a run. The rate constant kl was calculated from this equation, with the results given in Table 11; kn was g /1. Temp., Iotiic kl x lo3, Run " C strength [OCl-] [SO,-] [OH-] (g-1110l./l.)-~ III~II-I taken from reference 1, the following values being used: The values in Table I1 show that the rate is dependent on the hydroxide concentration; and, in fact, [OH-]. kl is approximatel!- constant, as shown by the following values, all for 60" C: Runs 1, 2 5, 6 7, 8 [OH-I. kl X lo

4 1648 CANADIAN JOURNAL OF CHEMISTRY. VOL. 39, 1961 This is equivalent to saying that the rate is proportional to the hydrogen ion concentration, and the most probable explanation is that this is really a reaction of hypocl~lorous acid, with rate = k[hocl]. [NO?-]. Nitrous acid is a much stronger acid than hypochlorous, so it is much more likely that the H+ is attached to the hypochlorite. On this assumptioil, the true rate constant, k, is given by where Kh is the equilibrium constant for the hydrolysis of hypochlorite ions: OC1- + H?O = HOCl + OH-. The ionization constant of hypochlorous acid is 3.8X10-8 at 27O C (6), and combining this with the ionization constant of water, Kh is 3.1 X10-7 at 27' C. At 25" C, AH for this hydrolysis is 11.0 ltcal; if we assume that this value holds up to 60 C, then Kh is 1.92X10-6 at GO0 C. This value cannot be claimed to be very accurate, but is the best available on the present data. Hence k = 1.61 X lo3 (g-mol./l.)-l min-l at 60' C. The data in Table I1 make the apparent activation energy 17.4 kcal/g-mol. However, since the heat of hydrolysis of hypochlorite ions is 11.0 kcal, the true activation energy is 6.4 kcal. If k is given by the usual formula k = Ae-EIRT, then log A = This is a rather small value, but there are others as small (see e.g. ref. 7). Runs 7 and 8, in conjunction with the others, show that the rate constant increases with ionic strength, though the change is fairly small. The total reaction here is 2. REACTION OF HYPOCHLORITE ASD IODATE IONS NaOCl + NaI03 + 2XaOH -+ XaC1 + Na31H?IOo. It was checked, by analytical methods given below, that the number of gram-molecules of periodate formed was within 1/2% of the number of gram-molecules of iodate disappearing. This confirms the stoichiometry of the reaction, and also checks the analytical methods. In alkaline solution a precipitate of sodium periodate forms. This was analyzed and found to be Na3H2IO6, in agreement with the phase diagram for the system sodium hydroxide, periodic acid, water, as found by Hill (8). The reaction was followed by mixing solutions of sodium hypochlorite and reagent grade sodium iodate and sodium hydroxide, with sodiunl chloride added to keep the ionic strength constant. The mixture was kept at constant temperature, and samples for analysis were taken at intervals. It was found that both hypochlorite and periodate ions absorb light strongly at about 300 mp, so spectrophotometric analysis was not possible. The following titrations were therefore used. JJethods of Analysis A sample was taken, cooled in ice for a short time, and any sodium periodate was filtered off. The solution was then allowed to warm up to rooln temperature, and aliquots were pipetted out for analysis. This procedure was necessary, as, otherwise, sodiulll periodate crystallized from the solution after pipetting. The aliquots were analyzed by the following methods.

5 LISTER AND ROSENBLUM: OSID.4TIOZi OF NITRITE AND IODATE IONS 1649 (a) The aliquot was diluted and inixed with a known volume of standardized arsenious oxide solution and a large excess of sodium bicarbonate. Then 3 g of potassiun~ iodide was added, and the excess arsenious oxide was titrated with standard iodine solution. In this titration the hypochlorite first oxidizes the arsenite to arsenate; and then, on adding potassiunl iodide, the periodate liberates iodine (which reacts with the arsenite) : Na3Hz ICI + Hz0 -+ NaI03 + 2KOH + 2NaOH + I?. Hence this titration measures hypochlorite plus periodate. It is essential to allow the sodium hypochlorite to react before adding potassiunl iodide. (b) The aliquot was mixed with excess dilute (about 0.7 N) sulphuric acid and boiled for 10 minutes. This treatment destroys hypochlorite, but did not affect the iodate and periodate present. The solution was then cooled, excess potassium iodide was added, and the liberated iodine was titrated with sodium thiosulphate. This titration measures iodate plus periodate. (c) The aliquot was boiled with dilute sulphuric acid as in (b). Borax was added in considerable excess (the ph being about 8); then potassium iodide was added and the liberated iodine was titrated with arsenious oxide. This measures the periodate alone. Froin these titrations the concentrations of hypochlorite and iodate ions in the solution could be calculated. The coilstancy found for the difference, [OCl-] - [IOc], besides being evidence of the stoichiometry of the reaction, is a check on the reliability of the analytical methods. In this reaction, chlorate formation accounted for less than 0.1% of the total decomposition. The results show that the reaction is first order in hypochlorite and iodate ions; hence the equations that governed the reaction with nitrite apply here also. The results of a typical run are given in Table 111. TABLE I11 (Run at 60.0' C; ionic strength 1.23; [NaOH] = M) [IOo-I, AT Difference log IOCl-l/[IO3-1 The values of the rate constant, kl, are given in Table IV. Inspection of these values shows that there is little change of the rate constant with hydroxide concentration. Apart from a degree of scattering in the constants, there seems to be a definite upward drift of the rate constant with hydroxide concentration. Roughly the constants at GO0 C obey the equation k1 = 5.08X10-~(l+O.l95[NaOI-I]~). The main reaction is therefore one of hypochlorite ions, presumably OC Cl , but there is some third-order reaction with hydroxide ions as well.

6 C.\N.\I)I.lN JOURXAL OF CHEMISTR\'. VOL Initial Temp., Ionic Final kl X 10?, Run " C strength [OCI-] [lo:-i [OH-] [OH-] (g-mol./l.)-1 11ii11-1 The dependence on temperature corresponds to an activation energy of 26.1 kcaljg-inol. Expressing kl as AAe-E'RT, log.-i = 15.8:3. This is a relatively high value, though not unprecedentedly high. 3. DISCUSSIOS OF RESlrLTS A number of oxidations by hypochlorite have now been examined by various worlcers, and it is of interest to see whether any generalizations can be made. Table V gives a.ac.ti\-ation Reactio~~ log energy, AHo, lxal kcal Rel. OCI- + I0:t HOCI + NO? OCI- + OCI OCI- + CIO? HOCl + OCX HOCI + Br OCI- + Mn0,p HOCl + HC.0, OCI- + CI0:)- Ver). slow OCI- +SO,-- 1'er)- last OC1- CX- 9, 9,,,,, OCI- + I /0Cl- - *Time in minutes. tthis rate is proportional to [fi;o-1'. summary of these reactions. The first colu~nn gives the reagents indicating whether hypochlorous acid or hypochlorite ions are involved. The activation energies and log -4 values refer to the reactions as written, allowing for hydrolysis to HOCl where necessary. The AH0 values also refer to the reactions as written, and assume that the reagents and products are all in aqueous solution. The last four entries are for reactions producing, respectively, perchlorate, sulphate, cyanate, and iodate ions. Here only qualitative observatiolls on the rates are available, but the rates are either very fast or very slow.

7 ~ LISTER AND ROSENBLL'M: OXIDATION OF NITRITE AND IODATE IOXS 1651 Table V perhaps chiefly emphasizes our limited kilowledge of the reactions of hypochlorites, but slight indications of trends are apparent. In the first place, log A generally falls when E does, so the actual rate constants are not as different as the activation energies would suggest. The fall in log --1 is not great enough to offset the change in activatioil energy; in fact it is about half as large as would be necessary for this. At present no general explanation is suggested. There is solme indication that, if the reaction is relatively exothermic, the activatioil energy is lower. This is perhaps to be expected, since if the new bond to oxygen is strong, this would tend to lower the energy of the activated coinplex as it is formed. However, this trend is not universal; for instance, the very exothermic oxidatioil of oxalate has quite a high activatioil energy. This particular reaction imay be a little different from the others, since it is appareiltly not an oxygen transfer but an electroil transfer; although it could be formulated either way. A more puzzling case is the reluctance of hypochlorite ions to oxidize chlorate ions to perchlorate, although this would be a fairly exotherillic process. Some of these reactions involve hypochlorous acid, and some hypochlorite ions, but at present no general explanation call be offered for this distinction. REFERESCES 1. M. W. LISTER. Can. J. Chem. 34, 465 (1956). 2. F. FOERSTER and P. DOLCH. Z. E1el;trochem. 23, 127 (1915). 3. L. FARKAS, M. LEWIN, and R. BLOCH. J..Am. Chem. Soc. 71, 1988 (1949). 4. M. W. LISTER and Y. YOSHINO. Can. 1. Chenl ( i. E..A. SHILOV. 1. ADDI. Chem. (u.s.s.r.) il9-14). d a. ~, 6. M. W. LISTER. Can. J. Chem. 30, 859 (1652j. 7. E. A. MOELWYN-HUGHES. The Itinetics of reactions in solution. 2nd ed. Oxford University Press. p A. E. HILL. J. Am. Chem. Soc. 50, 2678 (1928). 9. M. 11:. LISTER. Can. J. Che~n. 34, 489 (1956). 10. R. 0. GRIFFITH and A. McI<~owx. Trans. Faraday Soc. 28, 518 (1932).

+ 30C1- + H20 + 2HCO3- + 3C1- + Np.

+ 30C1- + H20 + 2HCO3- + 3C1- + Np. THE REACTION BETWEEN CYANATE AND HYPOCHLORITE' ABSTRACT The reaction between sodium hypochlorite and potassium cyanate in the presence of sodium hydroxide has been examined. The main products are chloride,