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 trans- mittance procedure. The same solu- tions 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-phenanthro- line method for fcrrous iron has been extensively studied by previous in- vestigators (2, 6, 9). Bccause interfer- ing ions were essentially absent in the samples used for the present investiga- tion, no additional work along these lines was undertaken. It is to be ex- pected that those ions which are them- selves colored or which form colored complexes with 1,lO-phenanthroline will

interfere. Some interferences can un- doubtedly be corrected by use of prop- erly 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 spectrophotom- etry, as presented, is believed to offer a new approach to the analysis of opaque or translucent solutions which may have numerous applications be- yond 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, Neo- cuproine, 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 Deter- mination of Traces of lletals,” 2nd ed., Interecience, Kevv York, 1950.

( 5) Snell, F. D., Snell, C. T., “Colori- metric 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 Analyt- ical Chemistry and Applied Spectroscopy, Pittsburgh, Pa., Februar 1956. Con- tribution 60, Research Zenter, Chem- strand Corp.

Color Complexes of Catechol with Molybdate

G. P. HAIGHT, Jr., and VASKEN PARAGAMIAN’

Chemistry Deparfmenf, Swarfhmore College, Swarfhmore, Pa.

A procedure for determining molyb- date 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 pH has been revealed. The dependence on pH suggests a reaction mechanism and structure for molybdate ion in the region of pH 7. A new complex con- taining equimolar molybdate and cote- chol 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.

OLORED complex compounds of C molybdate and catechol (0-dihy- droxybenzene) were prepared by Wein- land and coworkers from 1919 to 1926

1 Present address, Massachusetts In- stitute of Technology, Cambridge, Mass.

(4-6). Later workers have applied the color interaction in solution to the detection and determination of molyb- date. This paper presents a study of the equilibria involved in aqueous systems involving catechol and molyb- date.

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 to elucidate the formula of the complex. They reported working in basic solu- tions 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 to prevent oxidation of catechol (3) but to neutralize the base, forming a

sulfite-bisulfite buffer which stabilizes the p H near the neutral point. Com- plex formation has been found to 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 to 1 : 1 catechol-molybdate when the pH is changed from 6 to 2 or less. Equi- librium constants have been determined for the neutral and acid complexes. The equilibrium between the two com- pleses 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 spectrophotom- eter with a thermostated cell compart- ment maintained a t 26’ =k 0.5” C.

642 ANALYTICAL CHEMISTRY

2.0, A

WAVE LENGTH [ m r )

Figure 1 . Absorption spectra for molybdate- catechol complexes

A. lO-'M molybdate + 0.02M catechol in neutral woter E . B'. C. D.

1 O - W molybdate + 0.2M catechol in 0.1M HCl B corrected for absorption by free molybdate 1 O-3M molybdate in 1 M HCl 1OV3M molybdate in 1M HClO, 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 molyb- denum(V1) alone. This is character- istic 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 com- plexes 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 liters2 per mole2. The formation constant for the reaction l\lo(VI) + catechol k [Mo(catechol)] is 0.60/(H+) liters per mole, when hydrogen ion concentra- tion 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

W 0 Z 4 m [L

0 rT) m 4

1.6

1.2

0.8

0.4

MOLE PERCENT Mo!VI)

Figure 2. Continuous variation studies at con- stant catechol-molybdate concentrations

A. 5 X 1 O-3M (Mo" + catechol) in water, 3 8 0 mp E. 5 X 1 O-SM (MoJ" + catechol) in NazS206 + N a O H

(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 33% MoVI indicate 2 catechol:l MoV1 [neutral)- Peaks at 50% MoVI indicate 1 1 1 complex (acid)

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 in- tense color in neutral solutions is due primarily to the larger formation con- stant rather than to much greater absorption of light by the 2 to 1 com- plex, the molar absorbances being roughly the same order of magnitude.

In 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 com- plex. It is probable that between pH 2 and 7 equilibrium exists between the two complexes. However, the constant values obtained for equilibrium con- stants 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, in- dicate that these assumptions are valid. The equilibrium between the two com- plexes 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 ab- sorbances u:, az, , , , , for mixtures of the same initial concentration, c , of molybdenum(VI), and various con- centrations of catechol in large excess,

MOLARITY OF CATECHOL

Figure 3. Absorbance resulting from adding varying excesses of catechol to molybdate solutions

A. 10-4M molybdate In water. X = 3 8 0 mp B. 5 X lO-'M molybdate in sulflte-bisulfite

buffer. X = 4 5 0 m,u C. 5 X 1 O - W molybdate, (catechol) =

flve times scale 0.01M hydrochloric acid. X = 3 8 0 rnp

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 concentra- tion of uncomplexed molybdenum(V1). If M is obtained for one wave length, it is a simple matter to find i t 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 pH 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 con- sumption 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. Hydrox- ide ion destroys the complex by re- moving 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 molyb- denum(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 solu- tions. Complex formation then occurs by the reaction

OH

Alteriiatively, if the molybdate is a tetrahedral ion, it could become octa- hedral upon coordination with catechol without change in hydrogen ion con- centration as follows:

HO

dence of oxidation-reduction was ob- served in connection with the com- plexes in question. Molybdenum(V) failed to give color changes with catechol in solutions like those used

However, it seems more reasonable that the oxygen coordination number to molybdenum(V1) should not change in this reaction.

In base catechol becomes ( y o - .

It may be argued either that this species is unreactive or that in reacting ac- cording to the equation &MoOs-- + + 20H- the presence of hydroxide ion forces the equilibrium to the left, destroying the complex. Loss of protons by the postulated H,MoOs-- would only intensify this effect.

In acid the situation is complicated by polyacid formation by molybdenum (VI), which consumes acid and pre- sumably blocks one or two of the co- ordinating positions for the catechol. Alternatively the hydroxide ions co- ordinated to molybdenum(V1) might become coordinated water in acid,

2CaH4O2-H Q MOOz(O&aHa)2 4- 2H20

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 con- verted to MOO*++.

Further evidence that the coordina- tion 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 com- plex salts was difficult to remove, in- dicating a structure involving two hy- droxide and two oxide ions coordinated to the molybdenum(V1).

which would cause hydrogen ion to appear in the reaction products in place of water. In 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 com- plex 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 hydro- chloric acid than in perchloric acid.

Oxidation-Reduction. No evi-

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 poly- molybdate 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) ap- pear in aqueous solutions, for which molar absorbances and formation con- stants 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 bisul- fite and guard against acid-base ef- fects by carefully buffering the solu- tion a t pH 7 .

ACKNOWLEDGMENT

The authors acknowledge the sup- port 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, P., Arch. Pharm. 262,329 (1924).

RECEIVED for review March 8, 1958. Accepted January 25, 1960.

644 ANALYTICAL CHEMISTRY