Some Analytical Applications of Reaction-Rate-Promoting Effects The Tris( 1,l 0-phenanthroline)iron(ll)-chromium(v1) Indicator Reaction V. V. S. Eswara Dutt and Horacio A. Mottola Department of Chemistry, Oklahoma Sfate University, Stillwater, Okla. 74074
Several rate-accelerating effects on the oxidation of ferroin by Cr(VI) in sulfuric acid medium are reported. In all cases, these effects are observed in the earlier portions of the reaction profile and no catalyic cycle appears to be associated with them. Conversion of the rate-modifying species to an inactive form by destruction or, mainly, by complexation with Cr(lll) seems to account for the lack of catalytic cycle. Judicious choice of reaction conditions allows the determination (by initial rate measurements) of microgram amounts/milliliter of oxalic acid, citric acid, vanadium(lV), arsenic(lll), chromium(Vl), hexacyanoferrate(lll), and mg/ml of molybdenumWl).
Broadly speaking, kinetic methods of analysis are based on monitoring the rate of change in a reactant (or product) concentration. In catalytically-based methods, the main reaction in which the monitored species participates, written as a stoichiometric and over-all reaction, is termed the indicator reaction. Tabulation of some indicator reactions can be found in the literature (1-3). With only a few exceptions, these tabulations list reactions used for the determination of metal ions or inorganic anions that act as catalysts of the indicator reaction. Application of the modifying effect (inhibition) of complexing agents on the catalytic effect of metal ions adds a new dimension to the term indicator reaction and provides a useful probe for the end-point detection in some complexometric and precipitation equilibrium titrations (4, 5 ) . The present paper reports on some analytical applications using the oxidation of tris( 1,10-phenanthro1ine)iron(11), ferroin, by chromium(V1) in sulfuric acid medium. Ferroin is a well known oxidation-reduction indicator (6) and its oxidation to the iron(1II) complex, ferriin, can be rendered slow enough, by adjustment of the sulfuric acid concentration, to serve as an indicator reaction. Several chemical species are capable of accelerating the early stages of this oxidation. We have observed the following species to behave in this way: antimony(III), vanadium(IV), arsenic(III), molybdenum(VI), hexacyanoferrate(III), oxalic acid, and citric acid. Evaluation of these rate-modifying effects led to the development of simple methods for determining some of these species, as well as chromium(V1) in low concentrations and, in a rapid and simple manner by monitoring the rate of absorbance change due to conversion of ferroin to ferriin a t 510 nm. Oxalic acid, for instance, can be conveniently determined ( 1 ) K. B. Yatsimirskii, "Kinetic Methods of Analysis," Pergamon Press. Oxford,1966. pp 64-66. ( 2 ) 2. Gregorowicz and T. Suwinska, Chem. Anal. (Warsaw), 11, 3 (1966). (3) A . M . Gary and J . P. Schwing, Bull. SOC.Chim. Fr., 1972, 3657. (4) H . Weisz and S . Pantel, Anal. Chem. Acta, 62, 361 (1972) (5) H . A . Mottola, Talanta, 16, 1267 (1969). (6) A. A . Schilt, "Analytical Applications of 1 ,lo-Phenanthroiine and Related Compounds." Pergamon Press, Oxford, 1969, p p 103-1 1 2 .
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in human blood serum and urine by application of the ferroin-Cr(V1) indicator reaction and an initial reaction rate approach (7). The transient nature of the rate-modifying effect would indicate that the substance acting as modifier is concurrently either destroyed or rendered inactive. This type of behavior seems better regarded as promotion (8). This description results from adopting as definition of a catalyst, a chemical species which remains unaltered a t the end of each reaction cycle. Induced reactions (9, 10) can be considered as due to promoting effects. Not all promoting effects, however, seem to satisfy one of the practical criteria for induced reactions: "If a substance is slowly oxidized or reduced by a reagent, the velocity of such a process is increased during a concurrent rapid reaction between the reagent and another substance" (10). The vanadium(IV) and arsenic(EI) effects reported here, for instance, fit equally well both the induced-reaction requirements and the definition of promoting effect. The effects of oxalic acid, citric acid, hexacyanoferrate(III), and Mo(VI), however, do not satisfy the criteria listed above and do not qualify as induced-reaction effects. The adoption of the "promotion" concept is more general and makes a clear distinction from catalysis in mechanistic terms. Moreover, it is very probable that all induced reactions fit the mechanistic description of "promotion."
EXPERIMENTAL Apparatus. Equipment and technique to record absorbance us. time plots have been previously described ( 7 ) . Determinations in 1-cm cells were carried out, the reaction being initiated by injection of Cr(V1) solution, or of ferroin solution if Cr(V1) was to be determined. Temperature control was maintained with a Lauda/ Brinkmann Model K-2/R circulator and the use of a thermostatable cell adapter (Cary Instruments). Measurement of radioactivity was performed with a Packard Tri-Carb Liquid Scintillation Spectrometer, Model 3320. Reagents. Unless otherwise specified, reagents used were analytical grade or have been previously described ( 7 ) . Vanadium(IV) solutions were prepared from vanadyl sulfate reagent (Fisher Scientific Co.) purified by recrystallizing thrice from dilute sulfuric acid. Solutions of purified vanadium(IV) were standardized against cerium(IV) sulfate, and prepared in 0.10M sulfuric acid to prevent oxidation to vanadium(V). The 14C-labeled oxalic and citric acids were supplied by Amersham/Searle Corporation (Arlington Heights, Ill.). A mixture of PPO (Nuclear Associates, Inc., Westbury, N.Y.) and naphthalene in 1,4-dioxane was used as the scintillation solution ( 2 1 ) . Procedures. Experimental details for the detection and determination of oxalic acid have been given ( 7 ) . Similar conditions apply to citric acid detection and determination. A chart speed of 3 inchjmin and separate recording of absorbance us time curves is recommended. V. V. S. Eswara Dutt and H . A . Mottola, Biochem. Med., 9, 148 (1974). A . E. Marteli, Pure Appl. Chem., 17, 129 (1968). H . A . Laitinen, "Chemical Analysis." McGraw-Hill Book Company, Inc., New Y o r k , N . Y . , 1960, pp 462-471. I . M . Kolthoff and V. A . Stenger, "Volumetric Analysis," Vol. I. Interscience Publishers, Inc., New Y o r k , N . Y . , 1942, p 174. J . B. Birks, "An Introduction to Liquid Scintillation Counting," KochLight Laboratories Ltd., p 18.
Detection of Vanadium(IV) and Arsenic(II1). A 0.05-ml portion of the test solution is mixed with 0.05 ml of 1.0 X 10- 3M ferroin and 0.05 ml of 5M sulfuric acid in the depression of a spot plate. Distilled water (0.3 ml) and 0.05 ml of 1.7 X 10- 3M potassium dichromate are added and the solution is thoroughly mixed with a glass rod. Decolorization takes place in about 15 seconds in the presence of 1.9 pg of vanadium, or 4 pg of arsenic. Instantaneous decolorization is observed in the presence of 4 pg of vanadium or 15 pg of arsenic. A blank under identical conditions takes more than 5 minutes for complete decolorization. Determination of Vanadium(IV) and Arsenic(II1). Samples containing a t least 3 pg/ml of vanadium or 5 pg/ml of arsenic and ranging from 0.1 to 0.6 ml are introduced into a 1-cm glass cell containing 0.18 ml of 1.40 X 10- 3M ferroin, 0.18 ml of 5M sulfuric acid, and water to adjust to a final volume of 3.0 ml. The reaction is initiated by injecting 0.18 ml [for V(IV)] or 0.25 ml [for As(III)] of 1.7 X 10-3M potassium dichromate and the decrease in absorbance of ferroin is followed for 1-2 minutes. The concentrations of the sought-for species are evaluated from the initial rates .with reference to a working curve. It should be noted that addition of
Table I. E f f e c t s of F o r e i g n Species on the D e t e r m i n a t i o n of Chromium (VI), A r s e n i c ( I I I ) , and V a n a d i u m (IV)
Cr(V1) must invariably follow the addition of ferroin, since V(IV) and As(III), if allowed to react with Cr(V1) before addition offerroin, are easily oxidized and no accelerating effect can be observed. Detection of Chromium(V1). For the detection of chromium(VI), 0.1 ml of the test solution containing a t least 0.6 pg of chromium is mixed with 0.05 ml of 5M sulfuric acid, 0.05 ml of 0.01M oxalic acid, 0.2 ml of 1.0 x 10-4M ferroin and 0.1 ml of water. Decolorization of ferroin takes place in 5 to 10 seconds in the presence of chromium(V1). Determination of Chromium(V1). Samples containing a t least 1.7 pg of chromium/ml (0.3 to 1.8 ml in volume) are mixed with 0.3 ml of 0.01M oxalic acid, 0.15 ml of 5M sulfuric acid and sufficient water to bring the final volume to 3 ml. The reaction is started by injecting 0.15 ml of ferroin solution of concentration around 1.4 X 10- 3M. The decrease in absorbance is followed for 1 to 2 minutes and the concentration of Cr(V1) estimated from the initial rates and reference to a calibration curve (7). Interferences. Interferences in the determination of oxalic acid have been discussed earlier (7) with the exception of gluconic acid, which was found to cause no interference a t levels as high as 0.5 mg. Interferences for citric acid are like those for oxalic acid. The effects of several foreign species on the determination of chromium, vanadium, and arsenic(III) appear in Table I.
Maximum level not causing interference, in pgG
Foreign species
Silver (I) Copper (11) Nickel (11) Lead(I1) Cobalt (11) Zinc Cadmium Mercury(I1) Manganese(I1) Calcium Magnesium Aluminum Chromium(II1) Iron(II1) Vanadium (V) Molybdenum (VI) T u n g s t e n (VI) Fluoride Chloride Nitrate Acetate Phosphate EDTA Bi(II1)
215 250 200 20 120 130 225 200 2,b 15OC 320 100 108 200 20,b 5 5 0 " 10 50 50,
RESULTS AND DISCUSSION The chromium(V1) oxidation of ferroin in sulfuric acid medium can be formulated as in Equation 1, 3[Fe(l,10-phen)3]3i
+ Cr3+ + SO,'- + 3 HzO
(1)
if the analytical concentration of chromium(V1) is less than 4 x 10-4M (12) and no protonation of CrS0T2- occurs at [H+] 7 3.OM (13). Equation 1 chemically defines the indicator reaction on which this paper is centered. The general experimental rate expression can be written as: d [ferroin]
-
15"
190 200 186 500 100 185 100
a Determination according to the procedure in the text in the presence of 0 . 2 5 pg of chromium(VI), or 0 . 6 fig of vanadium, or 3 pg of arsenic. Tolerance limit for chromium(V1) determination. Tolerance limit for vanadium(IV) and arsenic(II1) determinations. Note. Sb(1II) interferes at all concentrations and could be determined in the microgram range (3 to 15 pg/ml) by a procedure similar to that described for As(II1). The need of high acid concentrations to keep Sb(II1) i n solution demands careful blank measurements and subtraction.
dt
=
k [ferroin] [Cr(VI)] [H2S0J2
(2)
Equation 2 and all other rate expressions in this paper have been formulated from individual order dependence studies. Temperature effect on k and all other rate proportionality constants reported in this paper is summarized in Table 11. The total reaction rate in presence of rate modifiers will include other terms on the right-hand side of Expression 2, and to increase the sensitivity of the determination of modifier, the rate contribution of the term shown for 2 must be minimized. Increasing concentrations (12) J. Y . Tong and E. L. King, J . Amer. Chem. Soc., 75, 6180 (1953). (13) G . P. Haight, Jr., D. C . Richardson, and N. H. Coburn, Inorg. Chern., 3, 1777 (1964)
Tab le 11. Temperature Effect on Rate P r o p o r t i o n a l i t y C o n s t a m t w ProDortionalitv constant k Reaction
Unmodified ( k M-3 m i n -1) I n t h e presence of ( k + M-5 min-1) I n t h e presence of (kf M-6 min-1) I n t h e presence of ( k ' M-4 min-1) In t h e presence of (k" M-4 min-1) a
10
oc
20
15 O C
2.68
x
103
3.27
x
oxalic acid
1.22
x
1013
1.49
x 1013
citric acid
1 . 6 0 X lo1*
vanadium(1V)
3.25
arsenic(II1)
x ...
109
103
... 3 . 5 8 x 109 1.24
x
109
oc
25 'C
4.00
x
103
4.80
1.93
x
1013
2 . 3 0 X 10'8
... 3.76
x ...
109
x
103
2.00
x 10'2
4.58
x
109
1.36
x
109
Above 25 O C , decolorization of ferroin is observed even in absence of chromium(V1).
A N A L Y T I C A L CHEMISTRY, VOL. 46, NO. 8, J U L Y 1974
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2
4
6
8
10
12
14
TIHEpl
Figure 1. First-order plot for the reaction in the presence of oxalic acid. Cr(VI): 8.33 X 10-5M, ferroin: 7.60 X 10-5M, sulfuric acid: 0.25M. (A) No oxalic acid added, ( 8 ) 1.00 X 10-5Moxalic acid, (C) 2.00 X 1 0 - 5 M oxalic acid, ( D ) 3.00 X 10-5M oxalic acid, ( E ) 4.00 X 10-5M oxalic acid
of sulfuric acid, as expected, progressively increase the rate of oxidation of ferroin, and the rate of the unmodified reaction can be accordingly controlled by adjustment of the sulfuric acid concentration. Hydrochloric acid can equally well be used as reaction medium but phosphoric acid does not allow good differentiation between unmodified and modified rates. The rate of reaction in presence of modifiers increases with chromium(V1) concentration but so does the rate of Reaction 1 and, in general, beyond 8.0 X 10-5M Cr(V1) the difference between the two rates is of little analytical utility. Reaction in the Presence of Oxalic a n d Citric Acid. In the presence of oxalic acid and citric acid, the following is obeyed:
-
d [ferroin] dt k + [ferroinI2 [Cr(VI)] [H2S0J2[Carboxylic acid]
(3)
In both cases the accelerating effect of carboxylic acid is exerted only temporarily. This can be seen in Figure 1, a first-order plot for the reaction in the presence of oxalic acid. The earlier part of the reaction, showing deviation from first-order behavior with respect to ferroin, delimits the region in which the carboxylic acid is affecting the rate of the reaction. Once the effect of the oxalic acid is exhausted, all over-all rates become equal to that in the absence of such acid. A similar behavior was observed with citric acid. Evidently, the acids are either decarboxylated or rendered inactive by complexation with some ionic species of chromium in the medium. Extraction with tri-n-butyl phosphate (14) followed by esterification with diazomethane and gas chromatography (15) failed to show any oxalic acid 4-5 minutes after the reagents had been mixed. However, the possibility that complex formation with Cr(II1) (16) prevented the extraction of oxalic acid into tri-n-butyl phosphate casts some doubts on the significance of the GC experiments. The use of 14C-labeled oxalic and citric acids showed that 10 to 15% of the acid is decarboxylated during the reaction. Separate experiments proved that this destruction is not due to the reaction of (14) P. M . Zarembski and A. Hodgkinson, Biochem. J., 96, 717 (1965). (15) G . Charransol and P.Desgrez, J. Chromatogr., 48, 530 (1970). (16) A. Werner and H. Surber, Justus Liebigs’ Ann. Chem., 406, 304 (1914).
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oxalic acid with chromium(V1) or with ferriin. Further, the chromium(III)-oxalate complex [prepared from chromium(V1) and oxalic acid (IS)]showed no accelerating effect on the rate of the ferroin-Cr(V1) reaction. These observations lead to the conclusion that the absence of regeneration of the active species is partly due to decarboxylation but mainly to complexation of the carboxylic acid with chromium(III). The possible formation of free radicals could not be detected by the “initiation of mercury(1) chloride precipitation” technique (17 ) because of slow precipitation of ferroin as an ion-pair with the chloro-mercury(II) anion present in the solution. Furthermore, precipitation of a ferroin-dichnomate ion complex precluded carrying out the test at higher reagent concentrations. From an analytical viewpoint, the recommended reagent concentrations for the determination of oxalic acid are found ( 7 ) 0.25M sulfuric acid, 8.3 X 10- 5M chromium(V1) and 6.8 X 10-5M ferroin. These concentrations equally apply to the determination of citric acid. Utilizing the experimental procedure described for oxalic acid, citric acid can be determined if present in concentrations ranging between 9 and 77 pg/ml in the final solution with a sensitivity of 0.0100 absorbance unit ml-I min-I pg-I expressed as the slope of initial rate us. citric acid concentration (18).The range of oxalic acid concentration amenable to determination is 0.9 to 4.5 pg/ml in the reaction solution mixture with a sensitivity of 0.083 absorbance unit ml- min- pg- Oxalic acid determination at low concentrations is not only of importance in clinical work ( 7 ) but carries industrial importance, since in some reagents the tolerable levels of oxalic acid are very low (19). Determination of oxalic acid by means of the procedure presented here is easily applied to such samples as formic acid, acetic acid, benzoic acid, and others of industrial interest and offers a convenient alternative to turbidimetric procedures (19). Wavelength scanning in the visible and UV region of the spectrum failed to show spectral evidence of an absorbing intermediate. It shall be noted here that, under the experimental conditions used in this study, both oxalic and citric acid are only very slowly oxidized by Cr(V1). Hence if a concerted pair of reactions with Cr(V1) as common reactant is pictured, it would involve two slow reactions and fail to fulfill one of the practical rules of induced reactions (IO). In an attempt to elucidate whether the decarboxylation of oxalic acid could be due to reaction with ferriin some observations were made which can be summarized as follows. Reaction of oxalic acid with ferriin under experimental conditions similar to those reported for the determination of oxalic acid gave no change in absorbance a t least up to 1 hour after mixing. Addition of excess oxalic acid caused an increase in absorbance only after an induction period of 30 minutes. Ferroin was then rapidly formed and soon the initial absorbance value of the corresponding concentration was attained. After the modifying effect of oxalic acid is complete, the important chemical species expected to be present in the reaction mixture are: Cr(III), ferriin, and oxalic acid. Since ferriin is an iron(II1) complex of 1,lO-phenanthroline and oxalic acid forms a stronger complex with iron(1II) than 1,lO-phenanthroline does, it is quite likely, considering the relative values of analytical concentrations employed, that ferriin dissociates, iron(1II)-oxalate complex is formed, and 1,lO-phenanthroline is liberated. This (17) A. Y . Drummond and W. A. Waters, J. Chem. SoC., 1953, 2836; 1953,3119. (18) H. Kaiser, Pure Appl. Chem., 34, 56 (1973) (19) A. I . Ryaguzov. I . V. Aleksandrova, and I E. Cherepenina, Zbur. Anal. Khim., 27,2469 (1972)
c
1
Figure 2. by Cr(VI)
2
3
4 TVIE,min
5
E
Absorbance vs. time curves for the oxidation of ferroin in the presence of vanadium(lV)
C r ( V I ) : 1.00 X 10-4M, ferroin: 6.84 X 10-5M, sulfuric acid: 0.30M. ( A ) No vanadium added, ( E ) 5.89 X 10-6M vanadium(1V). ( C ) 1.18 X 1 0 - 5 M vanadium(lV), (0)1.77 X 1 0 - 5 M vanadiurn(lV), ( E ) 2.36 X 10-5M vanadium(lV), ( F )2.95 X 10-5M vanadium(1V)
1,lO-phenanthroline may combine with Cr(II1) or remain as such if the Cr(1II) is initially'aquated or complexed with excess oxalic acid. If this actually happens, it is easy to see why decarboxylation of oxalate by ferriin does not take place. However, the formation of ferroin after the induction period needs explanation. Since iron(II1)-oxalate is known to be highly photoreactive, it could undergo photolysis generating iron(II), which upon combination with the free 1,lO-phenanthroline would regenerate ferroin. It shall be mentioned that iron(I1) forms a stronger complex with 1,lO-phenanthroline than Cr(II1) does. Also 1,lO-phenanthroline is a stronger complexing agent for iron(I1) than oxalate is. No formation of ferroin is observed when the same tests are performed in the dark, providing support to the above speculations. The photolysis of iron(II1)oxalate is reported as a chain reaction (20), and both induction periods and dependence of the reaction velocity on the surface area of the container are common in chain reactions. Actually, when the above reaction is separately performed in a spectrophotometric cell and in a watch glass, a longer induction period is observed for the reaction in the watch glass. Reaction in the Presence of Vanadium(1V). The following expression was observed to represent the rate in the presence of vanadium(IV): d[ferroin] dt h ' [ferroin] [Cr(VI)] [HzS0,]2 [vanadium] (4) Figure 2 shows the change in absorbance with time for the reaction in the presence of vanadium, and depicts the typical reaction profile of the modifying effects reported in this paper. As in the case of oxalic and citric acids, the modifying effect vanishes a short time after initiation of the reaction, and the over-all rates parallel the rate in absence of vanadium. A first-order plot derived from the curves shown in Figure 2 shows the same general characteristics as exhibited by that for oxalic acid (Figure 1). For the determination of vanadium, the recommended experimental conditions call for a sulfuric acid concentration of about 0.3M, 1.0 X 10- 4M chromium(V1) and 6.8 X 10- 5M ferroin. Increase of sulfuric acid concentration from 0.1M to 0.3M significantly increases the rate in presence of vanadium. Further increase in acid concentration, or chromium(V1) concentrations larger than 1.0 x lO-4M, (20) C A Parker, Trans Faraday S O C 5, 0 , 1213 (1954)
2
1
3
TIME,min
Figure 3. Effect of oxalic acid concentration on absorbance vs. time curves for the oxidation of ferroin by Cr(VI) C r ( V I ) : 3.33 X 1 0 - 6 M , ferroin: 7.2 X 10-5M, sulfuric acid: 0.25M. ( A ) No oxalic acid added, ( E ) 1.0 X 1 0 - 4 M oxalic acid, ( C ) 5.0 X 10-4M oxalic acid, ( D ) 1.0 X 1 0 - 3 M oxalic acid, ( E ) 1.0 X 1 0 - * M oxalic acid, ( F ) 2.0 X 10-2Moxalic acid
do not much increase the difference in rate between the unmodified and the modified reactions. The concentration range for vanadium determination lies between 0.3 and 1.5 pg/ml in the reaction solution mixture with a sensitivity of 0.505 absorbance unit ml-l min-l pg-'. The standard deviation for five replicate determinations of 0.60 pg vanadium was found to be 0.013 with an average of 0.612 pg. As in the case of oxalic acid, no spectral evidence of an intermediate was observed with vanadium(1V). Since the reaction paths do not sustain a catalytic cycle and vanadium is rendered inactive when converted to the V oxidation state, the accelerating effect can be classified as a promoting or rate-modifying action of the species V02+ on the redox reaction between chromium(V1) and ferroin. However, the oxidation of vanadium(IV) to (V) by Cr(V1) is a fast reaction and the effect exerted by vanadium(IV) on the slow oxidation of ferroin by Cr(V1) can equally well be viewed as a concerted induced reaction (9, 10). Reaction in the Presence of Arsenic(II1). Kolthoff et al. (21) stated that "in 1N sulfuric acid, ferrous phenanthroline is not oxidized by chromium(VI), however, on addition of some arsenic(III), the oxidation is induced instantaneously." These authors believe that the unstable chromium(IV) produced initially oxidizes the indicator. The following rate expression was experimentally derived for the reaction in the presence of arsenic(II1):
- d [ferroin]
-
dt h " [ferroin] [Cr(VI)I [As(III)] [HzSO4I2 ( 5 ) Absorbance us. time curves and the corresponding firstorder plots in the presence of various concentrations of arsenic(II1) resemble the curves in the presence of oxalic acid and vanadium(IV), and indicate that the modifying effect again disappears' a short time after initiation of the (21) I . M . Kolthoff, E. 6. Sandell, E. J. Meehan, and S Bruckenstein. "Quantitative Chemical Analysis." 4 t h e d , T h e Macmillan Company, London, 1969, p 766.
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reaction. Increase of sulfuric acid concentration, however, does not improve the sensitivity of the method. The optimum range for arsenic determination falls between l and 6 pg/ml in the reaction solution mixture with a sensitivity of 0.084 absorbance unit ml- min-I p c The standard deviation for five replicate determinations of 4.00pg of arsenic was 0.12 with an average of 4.07. Reaction Characteristics in the Determination of Chromium(V1). In the determination of oxalic acid the Cr(V1) concentration in the reaction mixture is kept constant and the decrease in absorbance of ferroin is followed. The method for the determination of chromium is just the reverse of the above, in that oxalic acid concentration is kept high and constant. The experimental rate equation is similar to Equation 3. In the absence of oxalic acid, no reaction could be detected between ferroin and Cr(V1) for a t least up to 15 minutes after mixing. With introduction of oxalic acid, the rate increases approaching an almost limiting value a t 10-2M oxalic acid. The absorbance us. time curves (Figure 3) exhibit a rising trend after passing through a minimum, in the presence of relatively high concentrations of oxalic acid. This is likely due to the reaction between ferriin and oxalic acid regenerating ferroin, the monitored species. Separate experiments with ferriin [prepared by oxidizing ferroin with Ce(N)] and oxalic acid under similar experimental conditions confirmed the reduction of ferriin by oxalic acid. The rate of the reaction in the absence of oxalic acid was found to be negligible at constant chromium and ferroin concentrations in the acidity range 0.05 to 1M HzS04. In the presence of oxalic acid (1.0 x 10-3M), however, the initial rate increases with HzS04 concentration up to 0.25M, and then remains constant with further increase of acid concentration. The optimum range for determination is 0.15 to 1.0 pg of Cr(VI)/ml in the final solution with a sensitivity of 0.609 absorbance unit ml-I min-l pg-l. The standard deviation for 5 replicate determinations of 0.50 pg of Cr(VI)/ml was 0.008 with an average of 0.52. Effect of Some Other Species. Among other species examined, hexacyanoferrate(D1) in the range of 5 to 30 pg/ml, and molybdenum(V1) in the range of 0.1 to 2 mg/ml in the reaction solution mixture, exert some accelerating effect on the oxidation of ferroin by Cr(V1). These effects, however, do not seem to offer sufficient advantage in sensitivity and limit of detection to deserve further analytical consideration. Since there is no possibility of electron transfer between Cr(V1) and Mo(V1) or Fe(CN)& these effects cannot be viewed as induced r.eactions. First-order plots for the reaction in presence and absence of hexacyanoferrate(II1) also show the general trend of Figure 1, reinforcing the overall argument for a promoting effect. In the case of molybdate, if ferroin and molybdate are mixed first, an insoluble material is formed which is difficult to dissolve even by addition of acid. To observe the modifying effect, molybdate should be added as the final reacting species. Hexacyanoferrate(1II) also forms an insoluble salt with ferroin but this is readily soluble in acid, so that the order of addition is unimportant in this case.
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CONCLUSIONS Certain chemical species have been found to exert a transient accelerating effect on the chromium(V1) oxidation of ferroin in sulfuric acid. In most of the cases, it is possible to determine the amount of rate modifier in solution by initial rate measurements. The rate modifiers reported here can be divided into three categories: those that cannot undergo any redox reaction, under the given experimental conditions, with either Cr(V1) or ferroin, such as hexacyanoferrate(III) and molybdenum(V1) ; those that undergo oxidation with Cr(V1) but very slowly, such as oxalic and citric acids; those that are oxidized reasonably fast by Cr(V1). The modifiers in the third category exert an accelerating effect only during the time they are undergoing oxidation. In any event, all rate modifications reported in this paper appear due to the effective participation of intermediate valence states of chromium, forming complexes with the modifying species. The rate of oxidation of ferroin by Cr(V1) is only affected in the earlier portions of the reaction and the disappearance of the modified rate appears to be due mainly to complexation of the rate modifier with Cr(III). The advantages offered by the reported procedure for the determination of oxalic acid have been summarized in an earlier communication (7). Very few methods are available in the literature for the determination of vanadium in the tetravalent state. The procedure described in this paper is more sensitive than that using self color of vanadium(IV) (22) and less cumbersome than the formaldoxime method (23). Determination of arsenic is usually carried either by the arsenomolybolic acid procedure (24) or by the Gutzeit method (24). The former requires the conversion of arsenic(II1) to arsenic(V), and the latter needs reduction of arsenic(II1) to arsine. Although both methods offer higher sensitivity than that reported here, they involve elaborate sample preparation steps and longer determination times. Further, in the arsenomolybdic acid method, pH is very critical and phosphate seriously interferes, while in the Gutzeit method, careful control of hydrogen evolution is essential during the reduction; besides, strong interference from Cu, Ni, Co, and Hg is observed. None of these disadvantages are associated with the procedure for arsenic(II1) presented in this paper. In the Cr(V1) determination, the selectivity is comparable to that of the established diphenylcarbazide method (24) but offers better sensitivity than the latter. Ferroin is a more stable reagent than diphenylcarbazide. However, the methods for citric acid, ferricyanide, and molybdenum(V1) are less sensitive than the established procedures for these species and hence no recommendation is made. Received for review September 26, 1973. Accepted March 4, 1974. This work was supported by the National Science Foundation (Grants GP-28207 and GP-38822X). (22) R. Santini, Jr., J. F. Hazel, and W . M. McNabb, Ana/. Chim. Acta. 6, 368, (1952). (23) M . Tanaka, Mikrochim. Acta. 1954, 701 (24) G. Charlot. "Colorimetric Determination of Elements." Elsevier Publishing Co.. Amsterdam, 1964, pp 176, 177, and 229.
