presented in Table I. Accuracy was also checked indirectly by a comparison with a second procedure. Four samples of bright fiber were analyzed for soluble iron by the previously devcloped transmittance procedure. The same solutions were then delustered by addition of titanium dioxide and measured by the reflectance procedure. The results were in substantial agreement, as shown in Table 11. Sample H,which shows the largest percentage error, is actually outside the scope of the reflectance method, because of its low iron content. Precision was determined by making eight replicate determinations on a particular sample of fiber. The values given in Table 111 show that a t an iron concentration of 2 p.p.m., the standard deviation found was 1 0 . 2 p.p.m. Interference in the 1,lO-phenanthroline method for fcrrous iron has been extensively studied by previous investigators (2, 6, 9). Bccause interfering ions were essentially absent in the samples used for the present investigation, no additional work along these lines was undertaken. It is to be expected that those ions which are themselves colored or which form colored complexes with 1,lO-phenanthroline will
interfere. Some interferences can undoubtedly be corrected by use of properly compensated blanks, as described above for fiber samples of differing initial color. Certain interferences will manifest themselves by a change in shape of the spectrophotometric curves. I n such cases it may be possible to employ a technique similar to that described by Diehl and Smith (4) for the simultaneous determination of iron and copper, utilizing measurements a t two wave lengths. The technique described can probably be improved by various modifications and refinements. However, the basic concept of reflectance spectrophotometry, as presented, is believed to offer a new approach to the analysis of opaque or translucent solutions which may have numerous applications beyond the scope of the one discussed. LITERATURE CITED
(1) Bandemer, S. L., Pchaible, P. J., IND.ESG. CHEX, A s . 4 ~ .ED. 16, 317 11944’1.
(2j Brokn, E. G., Hayes, T. J., Anal. Chim. Acta 7,324 (1952). (3) Davis, N. F., Osborne, C. E., Jr., Nash, H. A,, ASAL. CHEX. 30, 2035 (1958).
(4) Diehl, Harvey, Smith, G. F., “The
Copper Reagents: Cuproine, Neocuproine, Bathocuproine,” pp. 45-8, G. Frederick Smith Chemical Co., Columbus, Ohio, 1958. (5) Fortune, W. B., RIellon, M. G., IND. ESG. CHEM.,ASAL. ED. 10, 60 (1938). ( 6 ) Hoffman, C., Schweitzer, T. R., Dalby, G., Ibzd., 12, 454 (1940). (7) Hummel, F. C., Willard, R. H., Ibid., 10, 13 (1938). (8) Jackson, S.H., Ibid., 10,302 (1938). (9) Maute, R. L., Owens, 11. L., Jr., Slate, J. L., ASAL. CHEX 27, 1614-16 (1955). (10) Sloss, 11,L., llellon, 11,G., Smith, G. F., ISD. ESG.CHEX, ASAL. ED. 14,931 (1942). (11) Petersen, R. E., ASAL. CHEW 25, 1337 (1953). (12) Pflaum, R. I., Popov, A. I., Anal. Chitit. Acta 13, 165 (1955). (13) Pringle, n’. J. S , .4nalyst 71, 491 (1946). ( 4) Sandell, E. G., “Colorimetric Determination of Traces of lletals,” 2nd ed., Interecience, Kevv York, 1950. ( 5) Snell, F. D., Snell, C. T., “Colorimetric Methods of Analysis,” Vol. IIA, Van Sostrand, Princeton, X, J., 1959. RECEIVED for review September 11, 1959. Accepted December 31, 1959. Seventh Annual Pittsburgh Conference on Analytical Chemistry and Applied Spectroscopy, Pittsburgh, Pa., Februar 1956. Contribution 60, Research Zenter, Chemstrand Corp.
Color Complexes of Catechol with Molybdate G. P. HAIGHT, Jr., and VASKEN PARAGAMIAN’ Chemistry Deparfmenf, Swarfhmore College, Swarfhmore, Pa.
A procedure for determining molybdate developed by Seifter and Novic, involving formation of a complex with catechol, did not work when sodium sulfite was substituted for sodium bisulfite. Conditions for formation of the complex have been restudied and the need for careful control of p H has been revealed. The dependence on p H suggests a reaction mechanism and structure for molybdate ion in the region of pH 7. A new complex containing equimolar molybdate and cotechol forms in acid solution. The spectra and formation constants for the two complexes have been studied a t acidities in which only one or the other complex is formed.
C
complex compounds of molybdate and catechol (0-dihydroxybenzene) were prepared by Weinland and coworkers from 1919 to 1926 OLORED
1 Present address, Massachusetts Institute of Technology, Cambridge, Mass.
642
ANALYTICAL CHEMISTRY
(4-6). Later workers have applied the color interaction in solution t o the detection and determination of molybdate. This paper presents a study of the equilibria involved in aqueous systems involving catechol and molybdate. McGowan and Brian (2) indicate that the color results from a complex containing two catechol molecules per molybdate ion. Seifter and Kovic (3) found conditions for quantitative color development but made no attempt t o elucidate the formula of the complex. They reported working in basic solutions stabilized by addition of sodium pyrosulfite to prevent air Oxidation of catechol. Studies in this laboratory shorn that sodium sulfite yields entirely different results, virtually no comples being formed under conditions which were otherwise the same as those of Seifter and Novic (3). It is now apparent that the role of the pyrosulfite is not only t o prevent oxidation of catechol (3) but t o neutralize the base, forming a
sulfite-bisulfite buffer which stabilizes the p H near the neutral point. Complex formation has been found t o be most complete in neutral solutions. It drops off very rapidly in both acid and base. X o color is observed in 0.1M sodium hydroxide. The nature of the complex formed changes from 2: 1 t o 1:1 catechol-molybdate when the pH is changed from 6 to 2 or less. Equilibrium constants have been determined for the neutral and acid complexes. The equilibrium between the two compleses was not studied. EXPERIMENTAL
Reagent grade chemicals were used without further purification. Sodium pyrosulfite (Nai3iOs) is also called sodium metabisulfite. Baker and Adamson sodium metabisulfite and Llallinckrodt sodium bisulfite (NaHSOa) were used. Measurements were made with a Beckman D U spectrophotometer with a thermostated cell compartment maintained a t 26’ =k 0.5” C.
A
2.0,
W 0 Z
4
m
1.6
1.2
[L
0 0.8 rT) m 4
0.4
WAVE LENGTH [ m r )
Figure 1 . Absorption catechol complexes
spectra
for
MOLE P E R C E N T Mo!VI) molybdate-
+ 0.02M catechol in neutral + 0.2M catechol in 0.1M HCl
E.
lO-'M molybdate 1 O - W molybdate
B'.
B corrected for absorption b y free molybdate
C. D.
1 O-3M molybdate in 1 M HCl 1OV3Mmolybdate in 1M HClO,
A.
Catechol does not absorb in this region.
A Beckman Model G pH-meter was used to measure pH. RESULTS AND
CONCLUSIONS
Absorption srectra for molybdate and catechol mixtures are shown in Figure 1 for the near-ultraviolet region, where a maximum was observed by Seifter and Xovic (3). Readings in the vicinity of 400 mp and higher can be attributed to the complex alone, since molybdate and catechol do not absorb light by themselves a t these wave lengths. Absorption curves obtained in acid show a definite maximum a t 350 m l after subtraction of the curve for molybdenum(V1) alone. This is characteristic of a 1 to 1 complex, as shown by the continuous variation study in Figure 2. Much stronger absorption with a maximum a t 400 mp is observed in neutral solution ( A , Figure 1). This is due to a complex of 2 catechol with 1 molybdate, as shown by Figure 2. The molar absorptivities for the complexes a t wave lengths of maximum absorbance have been calculated and found to be 4.35 X IOa and 5.56 X 108 for the 1 to 1 and 2 to 1 complexes, respectively. The equilibrium constant for the reaction hIo(V1) 2 catechol k [l\lo(cate~hol)~] is 4.1 It 1.5 X 104 liters2per mole2. The formation constant for the reaction l\lo(VI) catechol k [Mo(catechol)] is 0.60/(H+) liters per mole, when hydrogen ion concentration is "between 0.1 and 1.OM." The hydrogen ion dependence diminishes a t pH = 2, where the value of K is 12 liters per mole. Calculations for the 1 to 1 complex formed in acid were made on the amumption that the complex contained one catechol per molybdate ion and that the same complex was formed in each acid concentration
+
+
woter
Figure 2. Continuous variation studies a t constant catechol-molybdate concentrations
+ + ++
5 X 1 O-3M (Mo" catechol) in water, 3 8 0 mp 5 X 1 O-SM (MoJ" catechol) in NazS206 NaOH (31,450 m l C. 10-2M (MoVI catechol) In 10-2M HCI, 4 5 0 mp D. 5 X 10-2M (MoVI catechol) In 1 M HCI, 4 3 0 mu E. 3 X 10-2hi (MoVI catechol) in 1 M HCIO4,400 mp Peaks at 3 3 % MoVI indicate 2 catechol:l MoV1[neutral)Peaks at 50% MoVI indicate 1 1 1 complex (acid) A. E.
+
+
used. The complex is too weak to obtain quantitative formation in large excess of catechol. The 2 to 1 complex could be prepared quantitatively and the molar absorbance observed directly (Figure 3). It is notable that the much more intense color in neutral solutions is due primarily to the larger formation constant rather than to much greater absorption of light by the 2 to 1 complex, the molar absorbances being roughly the same order of magnitude. I n 0.01M hydrochloric acid the continuous variation study indicated that there might be a small amount of 2 to 1 complex, because the maximum, which was not sharp, deviated from the position expected for a 1 to 1 complex toward that expected for a 2 to 1 complex. It is probable that between p H 2 and 7 equilibrium exists between the two complexes. However, the constant values obtained for equilibrium constants based on the assumption of pure 1 to 1 complex in acid and pure 2 to 1 complex in neutral solution, as well as the constant value for wave length of maximum absorption for acidities of 0.01M hydrogen ion or more, indicate that these assumptions are valid. The equilibrium between the two complexes which is pH dependent was not studied because of the complicated nature of molybdates in acid solution. The molar absorbance, M ,for the 1 to 1 complex was calculated as follows. Assume from the continuous variation study that the formula is 1 to 1 and that equilibrium exists. Using absorbances u:, az, , , , , for mixtures of the same initial concentration, c, of molybdenum(VI), and various concentrations of catechol in large excess,
MOLARITY OF CATECHOL
Figure 3. Absorbance resulting from adding varying excesses of catechol to molybdate solutions B.
10-4M molybdate In water. X = 3 8 0 mp 5 X lO-'M molybdate in sulflte-bisulfite
C.
5 X
A.
buffer.
X = 4 5 0 m,u
molybdate, (catechol) = flve times scale 0.01M hydrochloric acid. X = 3 8 0 rnp
1O-W
CI, CZ,etc. >> c, and KO,= (complex)
(MoVI)(catechol) then
may be used, because a is proportional to the concentration of complex and M c - u is proportional to the concentration of uncomplexed molybdenum(V1). If M is obtained for one wave length, it is a simple matter to find it for any wave length using the value of K., calculated a t the first wave length. VOL. 32, NO. 6, MAY 1960
643
Change of p H during Reaction. When 0.1M catechol is titrated with 0.1M molybdenum(V1) in water, a gradual change of pH from 5.3 t o 7.5 is observed. This is nothing more than would be expected from adding the weak base molybdate ion to the weak acid catechol. Therefore, there is no pronounced production or consumption of hydrogen ion in this reaction. Catechol probably displaces hydroxide ion from the coordination sphere of molybdate but must give up hydrogen ion to do so, resulting in a neutral reaction. Effect of Hydroxide Ion. Hydroxide ion destroys the complex by removing catechol according to the reaction CsH,(OH)z
+ OH- % CsHd02H- + Hz0
Hydroxide ion in excess of the catechol concentration destroys the complex completely. Partial neutralization of catechol gave results in fair agreement with calculations made assuming that only CJL(OH)2 reacts with molybdate and that Ki for catechol is I .4 X 10-'0. Structures of Molybdenum(VI)Catechol Complexes. Because molybdenum(V1) forms octahedra with oxygen and crystalline molybdates usually contain 2 moles of water of crystallization per mole of salt, we propose that the niolybdenum(V1) species is hfoO2(OH)4-- in neutral solutions. Complex formation then occurs by the reaction
Alteriiatively, if the molybdate is a tetrahedral ion, it could become octahedral upon coordination with catechol without change in hydrogen ion concentration as follows:
dence of oxidation-reduction was observed in connection with the complexes in question. Molybdenum(V) failed to give color changes with catechol in solutions like those used
HO
However, it seems more reasonable that the oxygen coordination number to molybdenum(V1) should not change in this reaction. I n base catechol becomes
(yo-.
It may be argued either that this species is unreactive or that in reacting according to the equation &MoOs-2CaH4O2-H Q MOOz(O&aHa)2 4-2H20 2 0 H - the presence of hydroxide ion forces the equilibrium t o the left, destroying the complex. Loss of protons by the postulated H,MoOs-would only intensify this effect. I n acid the situation is complicated by polyacid formation by molybdenum (VI), which consumes acid and presumably blocks one or two of the coordinating positions for the catechol. Alternatively the hydroxide ions coordinated to molybdenum(V1) might become coordinated water in acid,
+
+
OH
in this study. Solutions of the l to l complex in acid tended to turn blue on long standing or heating, indicating a secondary reaction in which polymolybdate may be reduced slowly by catechol to polymolybdenum blue. The findings of this research require no modification of the procedure of Seifter and Novic (3) for analysis. I t does clarify the conditions required for complex formation between catechol and molybdate, elucidating the nature of the complexes formed and the role of each reagent employed. SUMMARY
Two different complexes between catechol and molybdenum(V1) appear in aqueous solutions, for which molar absorbances and formation constants are given. The most suitable for analysis is the 2 to 1 complex formed in neutral solutions. The analyst should use precautions against, air oxidation of catechol by adding bisulfite and guard against acid-base effects by carefully buffering the solution a t pH 7 . ACKNOWLEDGMENT
The nonexistence of a triscatechol molybdate would seem to indicate that two oxygens'are not labile and do not take part in complex formation and that these two oxygens are those which would remain if the molybdate were converted to MOO*++. Further evidence that the coordination is octahedral is found in the work of Fernandes (f), who described a series of salts containing catechol and molybdate as anions represented by the formulas [ M o ~ & ~ o ] - - and [Moo2$:]--, where Ph is C&aO2--, Further, the water in the 1 to 1 complex salts was difficult to remove, indicating a structure involving two hydroxide and two oxide ions coordinated t o the molybdenum(V1).
644
ANALYTICAL CHEMISTRY
which would cause hydrogen ion to appear in the reaction products in place of water. I n such case hydrogen ions would inhibit complex formation. Substitution of perchloric acid for hydrochloric acid has very little effect. The equilibrium constant was calculated to be 0.8 0.1 liter per mole of 1M perchloric acid, assuming the same complex was formed as in hydrochloric acid. The slightly higher value of the constant obtained in perchloric acid may be caused by competition between catechol and chloride ion for complexing sites in hydrochloric acid. Corrections for molybdenum(V1) absorbance are greater in the 320-mp region in hydrochloric acid than in perchloric acid. Oxidation-Reduction. No evi-
*
The authors acknowledge the support of the Office of Ordnance Research, U.S. Army, in carrying out' this study. LITERATURE CITED
(1) Fernandes, L., Guzz. chim, itul. 55, 424 (1925). (2) McGowan, J. C., Brian, P. W., Nature 159,373 (1947). (3) Seifter, Sam Novic, Betty, ANAL. CHEM.23,188 (1951). (4) Weinland R.. F., Babel, Adolf,
Gross, Karf, Mal, Hermann 2. anorg.
allgem. Chem. 150, 177 (1926). (5) Weinland, R. F., Gaisser, F., Ibid.,
108,231 (1919). (6) Weinland, R. F., Huthmann, Arch. Pharm. 262,329 (1924).
P.,
RECEIVEDfor review March 8, 1958. Accepted January 25, 1960.
