Similarly, the plasma water values determined by several investigators and tabulated by Davis et al. (.?) were shown to be quite consistent, the mean being 92.5 with extreme values Of and These Our ex-
perience that sera water is quantitatively adequate as a reference for the determination of blood ethanol.
94x.
(3) F. E. ~ ( 1953).
~Keith~K
i ~ and ~ Jack ~ , Kirk, ~
science, ~ 118,~276
,
RECEIVED for review April 28, 1967. Accepted June 12, 1967. Opinions o r conclusions contained in this report are those of the author and d o not necessarily reflect the views or endorsement of the Navy Department.
Kinetic Study of Cerium(1V)-Vanadium(IV) Titration Reaction G . A. Rechnitz and G . N. Rao Department of Chemistry, State Uniuersity of New York, Buffalo, N . Y . 14214 USEOF THE CERIUM(IV)-VANADIUM(IV) reaction for the estimation of vanadium(1V) was first studied by Furman (I) who reported that, while potentiometric titrations in sulfuric acid media could be satisfactorily carried out a t room temperature, the attainment of steady potentials near the end point requires considerable time. Elevated temperatures (50"-60" C) facilitate the rapid attainment of steady potentials and permit satisfactory titrations in sulfuric, hydrochloric, and perchloric acid media (2). More recently, the beneficial effects of orthophosphoric acid (3) and of acetic acid, for the titration with ferroin indicator (4), in producing rapid end point attainment, even at room temperature, have been demonstrated. These observations and empirical studies have led to the impression that the sluggishness of the titration is due to the slow reaction between cerium(1V) and vanadium(1V) at room temperature. However, we found the cerium(1Vbvanadium (IV) reaction to be quite rapid ( k N 1350 M-l sec-I), even a t 25" C in molar H2S04, and, therefore, undertook to study in detail the kinetics of the cerium(IVbvanadium(1V) reaction as well as the reversible vanadium(V)-ferroin reaction in a n effort t o identify the cause of sluggishness in the analytical titration and to provide information useful for the selection of optimum reaction conditions. EXPERIMENTAL
The experimental arrangement and the procedure for the kinetic runs were as described earlier (5). Stoppered 4-cm silica cells were employed so that cerium(1V) concentrations in the range of 2 t o 10 X 10-5M could be determined spectrophotometrically a t 350 mp (E in 1M H2SO4a t 25" iv_ 4000). Cerium, vanadium, and ferroin solutions were prepared from reagent grade materials and were standardized by conventional methods (5, 6). Cerium salts and ferroin were obtained from the G. F. Smith Chemical Co. Potentiometric titrations were performed using standard procedures. The reactions involving ferric and ferrous o-phenanthroline complexes were followed by monitoring the ab(1) N. H. Furrnan, J. Am. Cl7em. Soc., 50,1675 (1928). (2) H. H. Willard and P. Young, Zrid. Eng. Chem., 20, 972 (1928). (3) G. Gopala Rao and L. S. A. Dikshitulu, Tu/un/a,9, 289 (1962). (4) K. Sriramarn and G. Gopala Rao, Tulurzfa, 13. 1468 (1966). (5) G. A. Rechnitz, G. N. Rao, and G. P. Rao, ANAL.CHEM., 38,
1900 (1966). (6) G. F. Smith and F. P. Richter, "Phenanthroline and Substituted Phenanthroline Indicators," G. F. Smith Chemical Co., Columbus, Ohio, 1944. 1 192
rn
ANALYTICAL CHEMISTRY
sorbance at 510 mp (molar absorptivity of ferrous o-phenanthroline a t 510 m p N 11000). RESULTS AND DISCUSSION
Several kinetic experiments performed with different initial concentrations of cerium(1V) and vanadium(IV), followed the second order rate law In
6 (a - x ) a (b - x)
=
kt (a
- b)
where a and b are the initial concentrations of the reactants and x the Ce(II1) concentration a t time, t . The rate constant was found to be 1350 =t100 M-l sec-l in 1 M H 2 S 0 4at 25" C with initial concentration of Ce(1V) varying from 2.5 to 9.0 X 10-5M and of V(1V) from 4.0 to 8.0 X 10-jM. Further experiments, performed in presence of 0 to 2.5 x lO-3M cerium(II1) and 0 to 2 X 10-4M vanadium(V), employing 5 X 10-jM each of cerium(1V) and vanadium(1V) in one molar sulfuric acid, indicated that the products d o not have any appreciable effect on either the form of the rate law or the rate constant. The absorbance values of cerium(1V) were corrected for the contribution from vanadium(V) in these experiments. The second order rate constant decreases slightly with increasing sulfuric acid concentration. Addition of different amount of sodium perchlorate at constant sulfuric acid concentration also results in a n increase in the value of rate constant, but the effect of solvent species upon the kinetics of the primary redox reaction was not studied in detail, because of the many possible reactant and product species present in these media (7-9). The temperature dependence of the rate constant gave the following thermodynamic values in M H2S04: activation energy = 13.0 + 1 Kcal., AH* = 12.4 i 1 Kcal., AS* = 2e.u. In view of the fairly high rate constant observed for the cerium(1V)-vanadium(1V) reaction, the need for elevated temperatures in the potentiometric titration cannot be due to the sluggishness of the primary reaction. On the other hand, the sluggishness of the titration could be due to slow attainment of potentiometric equilibrium. This possibility was examined by carrying out the titration (platinum indicator electrode GS. saturated calomel reference electrode)
+
(7) L. A. Blatz, J . Phys. Cliem., 66, 160 (1962). (8) T,J. Hardwick and E. Robertson, Cur?.J . Chem., 29, 828 (1951). (9) J. S. Littler and W. A Waters, J . Clzem. SOC.,1959,4046.
in the reverse direction-i.e., the titration of vanadium(1V) into cerium(1V). As expected from the potentiometric reversibility of the cerium(1V)-cerium(II1) couple, steady potentials were obtained within a few seconds prior t o the end point. In the end point region, 3 to 4 minutes are required for potentiometric equilibrium and, after the end point, extensive waiting is necessary between addition of vanadium(1V) increments before steady potential readings are attained. This observation is in agreement with the findings of Smith and Bannick (IO) who reported that nearly ten minutes are required t o obtain equilibrium potential values in the vanadium (1V)-vanadium(V) system in sulfuric acid media. We further found that the titration reaction between cerium(1V) and vanadium(1V) could be carried out rapidly and accurately in sulfuric acid media using spectrophotometric monitoring of the absorbance of the reaction mixture in the 400 to 430 m p wavelength range. It is, thus, clear that the observed sluggishness ( I , 2) of the cerium(1V)-vanadium(1V) titration a t room temperature is not due to slowness of the primary oxidationreduction reaction but, rather, the slow attainment of potentiometric equilibrium of the vanadium(V)-vanadium(1V) couple a t the indicator electrode. These results d o not, however, explain the sluggishness of this titration when carried out using ferroin as a visual indicator. Since the primary reaction is rapid, the origin of this effect is likely to be found in the interaction of one o r both of the primary reactants with the indicator. Hence, the interaction of the vanadium couple with the indicator was investigated in detail. The rate of oxidation of vanadium(1V) by the ferric o-phenanthroline complex [keeping the concentration of vanadium(1V) in excess to make the reaction pseudofirst order] obeys first order kinetics to about 7 0 x completion. The second order rate constant was found to be 5.7 + 0.1M-l sec-' in the concentration range of Fe(Phen)JII = 0.23 to 0.76 X 10-4M and vanadium(1V) = 0.84 to 3.34 X lOv3M a t 25" C in 1M H2SO4. Accurate values of the final equilibrium constant for this reaction are difficult to obtain because both the ferric and ferrous o-phenanthroline complexes undergo some dissociation in acidic media. Table I shows that the rate constant for the reaction decreases appreciably with increasing sulfuric acid concentration. The addition of acetic acid, a t constant sulfuric acid concentration, to the reaction medium resulted in an increase of the reaction rate while sodium perchlorate decreases the rate. Kinetic measurements carried out a t varying temperatures over the range of 15" to 35" C yielded a n activation energy of 12.0 i 1 Kcal and a n entropy change of -17 e.u. The reverse reaction-Le., the oxidation of the ferrous ophenanthroline complex by vanadium(V)-was followed under pseudo-first order conditions by keeping vanadium(V) in excess. The reaction reaches an equilibrium and the data were treated using the equation for a first order reaction opposed by a second order reaction (5). From experiments carried out over the concentration ranges Fe(Phen)JI = 0.30 to 0.76 X 10-4M and vanadium(V) = 0.83 to 6.67 x 10-3M, the second order rate constant was found to be 2.45 f 0.15 M-' sec-' at 25" C in 1M HzSO~. The reaction rate increases with increasing sulfuric acid concentration and decreases with the addition of acetic acid a t constant sulfuric acid concentration (see Table 11). The activation energy for the reaction was found to be 7.0 + 1 Kcal. from measurements taken a t 15", 25", and 35" C in molar sulfuric acid solution; AS* was calculated as -35 e.u.
shows that the primary redox reaction has a specific rate approximately 100 times as great as the vanadium-indicator reaction. Since the interaction of cerium(1V) with the indicator is known (12) to be very rapid ( k = 1.6 X 105 M-1 sec-1 at 25" C in M HzSO,), it is clear that the sluggishness of the titration carried out with ferroin indicator arises frcm the kinetic limitations of the indicator interaction with the vanadium(V)/(IV) couple rather than through any slowness of the primary redox reaction. Further, the beneficial influence of acetic acid (4) on the titration is seen to be due to
(10) G. F. Smith and W. M. Bannick, Jr., Talunfa, 2, 348 (1959).
(11) G. Dulz and N. Sutin, I m r g . Clzem., 2,917 (1963).
+
Table I. Rate Constants for Vanadium(1V) Ferric o-Phenanthroline Reaction in Various Media (Initial V(1V) concn. = 3.34 X 10-3M, initial ferric o-phenanthroline = 0.76 X 10-4M, 25" C) Concn of Concn of sodium sulfuric perchlorate Concn of Rate constant (M) acetic acid ( M ) M-1 sec-1 acid ( M ) 0.33 ... ... 10.4 0.67 ... ... 7.4 . . ... 5.7 1 .00 1.50 ... ... 4.0 2.00 ... ... 2.9 1.00 ... ... 2.90 1 .00 ... ... 11.3b ... 4.4 0.33 0.15 0.33 0.30 ... 3.0 , . . 1.4 0.33 0.60 0.33 1 .OO ... 0.64 ... 0.47 0.33 1.50 1 .oo ... 1,20 12.1 ... 2.00 20.6 0.67 0.67 ... 2.50 25.1 a at 15" C. bat 35" C.
+
Table 11. Rate Constants of Vanadium(V) Ferrous o-Phenanthroline Reaction as Function of Solvent Composition V(V) = 6.67 X 10-3M, Fe(phen)s I1 = 7.6 X 1 0 - 5 M , all at 25OC" Rate const, Solvent M-lsec-l 0.33M HzS04 0.12 0.67M HzS04 0.67 1. H2S04 2.45 14.40 1.50M HzSOa 2 . OOM H2S04 29.60 1 .OOM HYS04 1.56a 3.45b 1 .00M Has04 1.45 1 .OOM H2S04 1.OOM acetic acid 0.95 1 . OOM H2S04 2 , OOM acetic acid a at 15" C . b a t 35" C.
+ +
Inspection of the rate constants obtained for the reaction cerium(1V)
+ vanadium(1V)
+
cerium(II1)
+ vanadium(V)
(2)
and
+ ferric (o-phenanthr~line)~Fki! vanadium(V) + ferrous (o-phenanthr~line)~(3)
vanadium(1V)
k2
VOL. 39, NO. 10, AUGUST 1967
1193
Table 111. Comparison of Kegfor Vanadium-Indicator Reaction from Potentiometric and Kinetic Measurements (All at 25’ C) HSOa Formal potentials, volts concn, M Vanadium couple Indicator couple Keq(potent) kl M-1 sec-1 kz M-’ sec-l Keq(kinetic) 0.33 0.936 1.087 -350 10.4 0.12 87 0.67 0.986 1.073 -30 7.4 0.67 11 1.00 1.008 1.06 -1.6 5.7 2.45 2.3 1.50 1.030 1.045 ~ 1 . 8 4.0 14.4 -0.3 2.00 1.056 1.03 4 . 4 2.9 29.6 NO.1
its effect on the relative rates of interaction of vanadium(1V) and vanadium(V) with the oxidized and reduced forms of the indicator, respectively. Based on our knowledge of the predominant species present in sulfuric acid media (7-12), the most likely path for reaction 2 is Ce(SO&
+ VO(S0,) .HzO F? Ce(SO4)2-
+ VO(S0,) .HzO+
(4)
with a n equilibrium constant of 3.14 X lo6, in molar sulfuric acid solution, calculated from the formal potentials of the redox couples involved. Since both the forward and back(12) W. A. Waters and J. S. Littler, in “Oxidation in Organic Chemistry,” K. Wiberg, Ed., Academic Press, New York, 1965, p. 186ff.
ward rate constants of reaction 3 were measured in this study and since the formal potentials of the vanadium and indicator couples as a function of sulfuric acid concentration are available in the literature (6, IO), the equilibrium constants for reaction 4 obtained in our kinetic study are compared in Table I11 with those calculated from potentiometric data. Because the difference in formal potentials of the two couples is so small, the moderate uncertainty in the measured potentials produces a n especially large error in the calculated Kpot values. For this reason, our Kkin values calculated from the rate constant ratio are thought to be more reliable for this system.
RECEIVED for review April 7, 1967. Accepted June 5 , 1967. An Alfred P. Sloan Fellowship was awarded the senior author.
Separation of Amino Acids by Gas Chromatography Using New FIuoro Derivatives Glenn E. Pollock Exobiology DiGision, Ames Research Center, NASA, Moffett Field, CaW
IN THE COURSE of our research along other lines, it became necessary t o investigate a gas chromatographic technique for qualitatively identifying amino acids. A review by Weinstein ( I ) has thoroughly covered this field.
94035 W
z
4
EXPERIMENTAL
Several N-trifluoroacetyl (N-TFA) amino acid n-butyl esters were synthesized ( 2 , 3). Identical procedures were used for the N-pentafluoropropionyl (N-PFP) and N-heptafluorobutyryl (N-HFB) derivatives. G a s chromatography was carried out on a Perkin-Elmer 800 instrument, using Carbowax 20M, 0.02-inch X 150-foot columns and a linear pressure programmer. RESULTS AND DISCUSSION
The retention times of the high-boiling esters were long, resulting in analysis times well over 1 hour, while the retention time of the N-pentafluoropropionyl (N-PFP) derivative was significantly shorter. Because of the pronounced effect of N-PFP derivatives, we prepared N-heptafluorobutyryl (N(1) Boris Weinstein, “Methods of Biochemical Analysis,” Wiley, 1966, p. 203. (2) W. M. Lamkin and C. W. Gehrke, ANAL.CHEM., 37, 383 (1965). (3) I. Halasz and K. Bunnig, Z . Anal. Chern., 211, 1 (1965).
1 194
ANALYTICAL CHEMISTRY
Figure 1. Separation of the N-TFA, n-butyl esters of 14 amino acids Carbowax 20M, 0.02-inch X 150-ft column Temperature: Isothermal at 100” C for 2.5 min. Raise to 170” C at 4 “ C/min Gas: Helium, 9.5 Ib (about 10 mlimin) for 25 min. Then raise linearly to 20 Ib (28 mljmin) at rate of 7.74 Ib/min. These conditions are identical for Figures 1 and 2. Figure 3 had no carrier gas flow rate increase
HFB) derivatives and found that their retention times were even shorter. In Table I are shown the retention times, relative retention times, and per cent time reduction for each amino acid under isothermal conditions using a Carbowax 20M column. Table I shows that the N-PFP and N-HFB derivatives have
