George H. Schenk Wayne State University Detroit, Michigan
The Use of Inorganic Redox Mechanisms in Analytical Chemistry
M o s t elementary analytical textbooks have overemphasized the usefulness of standard reduction potentials and the Kernst equation for predicting whether redox reactions will occur quantitatively; i.e., whether a redox titration is feasible. For example, in the simple case involviug one electron in each half-cell reaction, it may be stated that the reaction is 99.9% complete if (Eo,.,. - Eorad .,.) 2 +0.354 volts. If the half-cell reaction involves hydrogen ions and the pH is not 0, then the half-cell potential, E, must be calculated using the Nernst equation. E is then used in place of a standard reduction potential in the comparison above. The result has been that much interesting chemistry has been deemphasized in favor of calculations of E and comparisons of En's and/or E's. I n addition these comparisons are not always strictly applicable. For instance, although the formal potential (sulfuric acid) of the Ce(1V)-Ce(II1) couple is 0.91 volts more positive than the Eoof the As(V)-As(II1) couple, the reaction is so slow in 1 M sulfuric acid that no reaction appears to occur. However, the presence of a trace of iodide ion catalyzes the reaction so that equilibrium is reached rapidly. An understanding of the catalysis of the reaction is therefore more important than a prediction about the point of equilibrium. It is also probably true that most textbooks have underemphasized the utility of redox mechanisms and kinetics, both of which aid in understanding catalysis. Only recently Laitinen (16) and Duke (7, 8) have reviewed many inorganic redox mechanisms pertinent to analytical chemistry. Moody and Thomas (18) have indicated some experimental applications of the mechanisms of the oxidation of iodide ion, and Schenk (22)has indicated some experimental applications of the mechanism of the periodate cleavage of glycok. It is interesting to compare the information provided by a subtraction of E O or E values and that provided by redox mechanisms coupled with some kinetic information. The former gives quantitative information about the conversion of reactants to products, the effect of pH, and the effect of ligands; the latter gives only qualitative information. En must be known before interferences can he predicted; these interferences can only be predicted on a thermodynamic basis for reactants and products. A knowledge of mechanisms, however, makes it possible to predict interferences on a kinetic basis for intermediates as well. It is also possible to extend the scope of some Presented in part before the Division of Chemical Education a t the 144th Meeting of the American Chemical Society, Los Angeles, California, April, 1963. Supported in part by Public Health Research Grant RG-7760 from the National Institutes of Health, Public Health Service.
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methods to unknown compounds, to predict products and thence equivalent weights. Finally, mechanisms make it possible to predict pathways for catalysis and to design trace analysis methods based on this knowledge. Since Duke and Laitioen concentrate on mechanisms involving cations or strong oxidants, this paper will deal mainly with common anion-anion or anion-molecule reactions where Eoor E values are not always useful or are not known. These types of reactions should be readily appreciated by students who have already encountered second order nucleophilic displacement (SN2)reactions. Iodine-Arsenic(1ll) Reaction
A simple system for laboratory study a t the undergraduate level is the equilibrium-controlled iodinearsenic(II1) reaction.' This reaction may be experimentally studied a t four different pH's: 1,4.7, 7.5, and 13. It will be observed that the titration reaction (at the end point in particular) is rapid a t pH's 7.5 and 13, hut stoichiometric only a t pH's 4.7 and 7.5. The student should conclude that the rate of the reaction appears to be subject to specific base catalysis; too high a concentration of hydroxide ion, however, can alter the stoichiometry. Instructions for an illustration experiment follow. More complete details will be found in "Laboratory Manual of Chemical Analysis," published by the author, 1962.
Experimental Study of As(lll)-1%Reaction Bufe?, p H 4.7: Dilute 60 g of glacial acetic acid and 82 g of anhydrous sodium acetate to 300 ml with water. Arsenic (ZZI), 0.1011 N : Dissolve exactly 2.5000 g of primary standard anenic(II1) oxide in a freshly prepared solution of 5 g of sodium hydroxide pellets dissolved in 20 ml of water. Add 50 ml of water, 10 ml of 12 M hydrochloric acid, and 1 g of sodium bicarhonate. Dilute to 500 ml, checking the pH to be 7-8. Iodine, 0.lN: Dissolve about 12.7 g of iodine in a solution of 40 g of potassium iodide dissolved in 25 ml af water. Dilute to a liter. Sodium thiosulfate, 0.lN: Dissolve about 12.5 g of NalSzOs.5H10 in 500 ml of water that has been recently boiled and cooled. Add about 0.1 g of sodium carbonate and 0.1 g of EDTA as preservatives. Osidation of As(III) at uarious pH's: Titrate a 25 ml aliquot of ~nenie(I1I)with iodine after d d i n e starch and the aooronriate reagent tb adjust the pH. The endupoint is taken ;sLtbe first starch-iodine color which persists after 2 minutes of hand mixing. For pH 1, add 0.5 ml of 12 A4 hydrochloric acid. Far pH 4.7, add 4 ml of the acetic acid-sodium acetate buffer. For pH 13, add 0.6 g of sodium hydroxide pellets. For pH 7.5, add 3 to 4 g pf sodium bicarbonate. Note the rate of reaction near the end point a t each pH. Reve~sibililyof the ozidation: T o the flask buffered a t pH 7.5 add droowise 3 ml of 12 M hvdrochloric acid. swirlin. between additio& Test with pH pape;to check that thk solution is acidic. If not, add more acid. Then add 40 ml of 12 M hydrochloric acid, making the solution about 4 t o 5 M in acid. Titrate the iodine released with 0.1 N NanS201after adding 2 g of ICI. Compare meq of arsenic(V) m d (111).
The titration a t pH 4.7 is a challenge to the student's judgment since the reaction is found to be quantitative only after patient mixing and dropwise titration a t the end point. The conditions can be varied using a buffer such as potassium acid phthalate (pH 4.0). At this pH the reaction is slightly slower and usually will fall just short of being quantitative. Instead of conducting a titration at pH 13, the conditions can be varied by using ammonium hydroxide to give a pH of about 11.3. At this pH the volume of iodine reagent consumed is only about 0.4% higher than that a t 7.5. The reversibility of the iodine-arsenic(II1) reaction is also tested bv acidifvinp the solution a t wH 7.5 after " reaching the equivalence point, adding excess potassium iodide, and titrating the liberated iodine with sodium thiosulfate. A probable mechanism (if one exists) for the iodinearsenic(II1) reaction may be:
cause of the higher hydroxide ion concentration, House (14) indicates the reaction is first order in persulfate and follom a free radical mechanism. Iodine-Thiosulfate Reaction
In contrast to the equilibrium controlled arseniciodine reaction is the classic irreversible iodine-thiosulfate reaction, studied in detail by Awtrey and Connick (2, 3 ) :
u
Obviously the principal equilibrium forms of arsenic(111) and arsenic(V) will depend on the pH, as indicated by Laitinen (16). Students who know about SN2 reactions will recognize the arsenic(II1) anion as a nucleophile in Reaction (1). It can be pointed out that the empty 4d orbitals on arsenic(II1) can accommodate some of the electrons of the iodine molecule in the transition state, making arsenic(II1) a good nucleophile. This polarizahility property has been invoked by Edwards and Pearson (9) to account in general for the nucleophilicity of very weakly basic anions. It should be emphasized that iodide need not he displaced simultaneously with the attack of hydroxide in Reaction (2) since arsenic, like phosphorus, should be capable of forming a pentadenate intermediate in the transition state. This can then decompose to products a t a later stage. The effect of hydroxide on the rate of the reaction can be explained by either the nucleophilic attack of hydroxide or the hydrolytic attack of water on the intermediate, as well as perhaps by the removal of the strongly acidic forms of arsenic(V) from the equilibrium in Reaction (2). The mechanism might be used to rationalize the observation that arsenic(II1) sulfide cannot be readily oxidized to arsenic(V) in nitric acid for qualitative analysis purposes (13). Although the standard reduction potential of the HNOa-NO couple is 0.46 volts more positive than the standard reduction potential of the As(V)-As(II1) couple, the detection and estimation of amounts of arsenic(V) is apparently not practical by using this reaction. Assuming that arsenic(II1) sulfide is readily converted to soluble arsenic(II1) ion, one might deduce that the rate of oxidation of arsenic(II1) is limited bv the low hydroxide ion concentration of nitric acid. It is possible to oxidize arsenic(II1) from the sulfide by use of persulfate in basic solution. Although one might predict that this oxidation would be rapid be-
The standard potential for the tetrathionate-thiosulfate couple cannot strictly be used to predict equilibrium considerations because of the irreversible second step. For the same reason, the reaction rate cannot be altered by removing the product tetrathionate. Tetrathionate cannot be determined by reduction with iodide, a reaction suitable for determining arsenic(V). Similar to arsenic(II1) in the arsenic(II1)-iodine reaction, thiosulfate behaves as a nucleophile by displacing iodide ion, this time in both reaction steps. Since iodide is only a slightly poorer nucleophile than thiosulfate, it can displace the latter from "positive" iodine in the intermediate by coordinating with the "positive" iodine to produce molecular iodine. Apparently the nucleophilicity of iodide toward tetrathionate is markedly lowered by the change in substrate (Q), and it cannot displace thiosulfate from tetwthionate in a detectable amount. Thus Reaction (4) is irreversible. The nucleophilicity of thiosulfate and iodide can be rationalized by the ability of empty d orbitals on sulfur or iodide to accept electrons (9). The titrimetric results of Kolthoff (15) involving 0.001 M solutions of iodine and thiosulfate provide an instructive application of the mechanism. When 0.001 M iodine was titrated a t pH 4.0-5.0 the formation of hypoiodite removed a significant amount of the iodine from the equilibrium in Reaction (3). Since hypoiodite oxidized thiosulfate to sulfate, the results were 3 to 4y0 low in the volume of thiosulfate titrant consumed. As the pH was lowered toward zero and the iodide concentration was increased (favoring the formation of 13- rather than 01-), the results were only 1%low. Kolthoff (15) also observed the reappearance of iodine a t the end point in his experiments; this was explained by the observation of Awtrey and Connick (8) that a t iodide concentrations below 0.07 M the following occurred :
Apparently below 0.07 M iodide, the intermediate's concentration can build up as a result of Reaction (3). Then, near the end point the thiosulfate concentration becomes low enough to permit the intermediate to react with itself, probably by a concerted coordination of iodide, which removes "positive" iodine from the intermediate behaving as a nucleophile. That Reaction (5) requires a minimum concentration of iodide Volume 41, Number I, January 1964
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is verified by the observation of a measurable reaction above 0.007 Ii4 iodide (2). Below 0.003 A4 iodide, the intermediate is oxidized to sulfate by iodine or hypoiodite (15). Uses of the Iodine-Thiosulfate Mechanism
The mechanism in Reactions (3) and (4) may be extrapolated to predict that reagents that can coordinate with iodine (and undergo overall 1 electron changes) may react similarly and may interfere in the iodine-thiosulfate reaction. Table 1 gives some examples, such as mercaptans (as RS-), xanthate (from carbon disulfide), and thiourea. For each there is an overall 1 electron change and the products are dimers. Table 1 . Specien 2SS03-2 2RS(Mercaptan) 2EtOCSS2(NH2hCS: (N&).CS:
sSCN-a
----
Swecies Oxidized bv Iodine Equiv. wt. of sodium salt of speries
Product S,0rc2 RSSR
(Dimers)
iormula wt formula wt
(EtOCSS), CdlN4He(+fast
Table 2.
Na-
Table 1 also contains examples of sulfur-containing anions that do not react by this mechanism, such as sulfite and thiocyanate, which undergo overall 2 and 8 electron changes, respectively. An interesting experiment a t the graduate level is to present the student with two possible mechanisms for the reaction of iodine with thiocyanate: the mechanism similar to Reactions (3) and (4), to give (SCN)2, and the following mechanism (1): 01-
Once the OzN-SS03- intermediate is formed, it is undoubtedly subject to hydrolytic attack by water to rupture the sulfur-sulfur bond and to rapid oxidation by iodine to give sulfite which may be further oxidized to sulfate.
Cyanide is not the only uucleophile that can cleave the S S linkage. Sulfite (21) has been used to cleave dixanthogen, (EtOCSS)2, to thiosulfate, which can be titrated with iodine. Trace Analysis
Mechanisms involving anions have also been used in the design of trace analysis methods based on catalysis. The mercaptan catalysis of the iodine-azide reaction to produce nitrogen and iodide is well known and useful for determining cystine and cysteine (23). A novel reaction of thiosulfate has been reported by Burkhalter (5) for the trace analysis of copper in semiconductor materials. Salicylate ion is added to form the colored Fe(II1)-salicylate complex; then Cu(I1) is used to catalyze the measurably slow Fe(II1)-thiosulfate reaction:
Parts per hillion of Cu(I1) can be determined by plotting the rate of fading of the red iron(II1) salicylate complex against the concentration of Cu(I1). The gross mechanism for electron transfer and catalysis is suggested to be: Cu + Z
+ M + 2SS03-s +"
-
CU
Cu
+
+
+2+n-4
+ SIOs-z
Mi&-1
(11)
As shown in reaction ( I t ) , Cu(I1) may coordinate with an electron pair on sulfur along with Fe(III), or possibly with another Cu(I1). Whether electron transfer from sulfur to the half-filled 3d orbital of Cu(I1) occr~rsimmediately, or when tetrathionate is formed, is open to spcculation. However, it is suggested that this double electron bridge may facilitate the electron transfers as well as thc formation of tetrathionate. The mechanism permits t,he predict,ion of int,erfrrcnccs such as other ions containing divalent sulfur and cations which are reduced by Cu(1) or Fe(I1) as well as species which will competitively coordinate with t,hiosulfate. Literature Cited
ANGELESCU, E., AND POPESCU,1.. D., Z. Physzk. Chem., 156, 258 (1931). R. E.. J. A m . Chem. Soe., (21 AWTREY.A. D.. AND CONNICK.
(1)
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, I)., A R D CONNICK, R. E., Ihid., 4546. (3) A v r n ~ u A. J. K., AXD SKOOO, 1). A , Anal. Chem., 26, 1008 ( I ) HARTLETT, (lM4). T ES., R ,Anal. Chem., 33, 21A (May, 1361). (5) ~ ~ E R K H A LT. (6) Iloor,; G., A N D GRIFFITH,R. O., Trans. Paraday Soc., 45, 546 (1940). (7) I)UKE,F. R., Anal. Chem., 31, 527 (195Y). ( 8 ) I)UKE,P.R., "Treatise on Analytical Chemistry," Part I, \.J., AN,, THOMAS, J. 1). It., CIIEM.~ U C 40, . , 151 (18) MOODY, (1963). \-.,--,.
W., "Mechnni~rn~ of Sulfur Reactions," MeGrsw(20) PRYOR, Hill, Ine., New York, 1962, p. 64. MCTRTHY, A. R., (21) ~ A T Y A N A R A Y A RAO,V. R., A N D VASUDEVA Chemist-Analyst, 50, 30 (1961). G. H., J. CHEM.EDUC.,39, 32 (1962). (22) SCHENK, It. I)., MACK,P. A,, AND C H I ~ ~W. S , A,, (23) STRICKLANI), Anal. C h e m , 32, 430 (1960). E. A., J. Chem. Soe., 101, 2166 (1912). (24) WERNER,
