1232
I n d . Eng. Chem. Res. 1987, 26, 1232-1234
Ouchiyama, N.; Tanaka, T. Ind. Eng. Chem. Fundam. 1981,20,67. Satterfield, C. N. Mass Transfer in Heterogeneous Catalysis; MIT Press: Cambridge, MA, 1970. Smith, J. M. Chemical Engineering Kinetics, 2nd ed.;McGraw-Hill: New York, 1970. Smith, D. M.; Huizenga, D. G. Proceedings of the 10th IASTED Symposium on Applied Modelling and Simulation; ACTA Press:
New York, 1984; p 13. Strober, W.; Fink, A. J . Colloid Inter. Sci. 1968, 26, 62. Wright, T. M.S. Thesis, University of New Mexico, Albuquerque, 1985.
Receiued for reuiew March 25, 1986 Accepted March 19, 1987
Stoichiometry of Ferrous Nitrosyl Complexes David Littlejohn and S. G. Chang* Lawrence Berkeley Laboratory, MIS 7O-llOA, Applied Science Division, University of California at Berkeley, Berkeley, California 94720
The ferrous ion t o nitric oxide ratio has been determined for the class of compounds Fe"(L),(NO),, where L = H20, citrate, IDA, NTA, and EDTA. All compounds studied were found to have a maximum ferrous ion to nitric oxide ratio of 1:l. In addition, the equilibrium constant for the binding of nitric oxide to Fe"(H,O), was determined to be 470 M-l a t pH 3 and 25 O C . A number of ferrous chelates rapidly and reversibly react with nitric oxide in aqueous solutions (Littlejohn and Chang, 1982) and have been proposed as agents to aid in the removal of nitric oxide from power plant flue gas streams (Chang et al., 1982). The strength of the ferrous ion-nitric oxide bond is highly dependent on the chelate used in the complex (Griffiths and Chang, 1986). There has been some discrepancy in the literature about the stoichiometry of a number of ferrous nitrosyl complexes. McDonald et al. (1965) studied several ferrous nitrosyl complexes by electron spin resonance and found both mono- and dinitrosyl complexes, depending on the ligand involved. Hishinuma et al. (1979) measured the absorption of a diluted nitric oxide gas stream passed through a solution containing Fe"(EDTA), where EDTA = ethylenediaminetetraacetate, and found a mononitrosyl complex. Pearsall and Bonner (1982) report evidence for a dinitrosyl complex, as well as mononitrosyl complex, in ferrous ion solutions with an acetate buffer. Ogura and Ishikawa (1983) and Ogura and Watanabe (1982) investigated the complexation of nitric oxide with Fe"-aminocarboxylic acid complexes by electrochemical and spectroscopic means and found Fen-to-NO stoichiometries of 1:2 for the EDTA complex, 1:2 and 1:l for NTA (nitrilotriacetate)and IDA (iminodiacetate) complexes, and 1:l for EDDA (ethylenediaminediacetate) and glycine. Because of the conflicts about the stoichiometry in the literature, we decided to study a number of ferrous nitrosyl complexes that have potential for use in flue gas scrubbing systems. We have investigated the stoichiometry of the binding of nitric oxide to Fer1(H20),,Fe"(cit) (where cit = citrate), Fe"(IDA), Fe"(NTA), and Fe"(EDTA) by measuring the absorption of nitric oxide by solutions of these compounds. Experimental Section Nitric oxide (Matheson, C.P.) was purified by three trap-to-trap distillations in which the first and last portions were discarded each time. The purified nitric oxide was stored a t liquid nitrogen temperature until use. Reagent-grade chelates and ferrous ammonium sulfate were used to prepare the ferrous chelate solutions. We believe the ammonium and sulfate ions do not appreciably influence the chemistry of this system. Prior to addition of the ferrous salt, the chelate solutions were thoroughly degassed on a vacuum line by repeated evacuation and cavitation. The bulb containing the solution was filled with argon, and 0888-58S5/87/2626-1232$01.50/0
the ferrous salt was added. After addition of the ferrous salt, the solutions were treated with concentrated HC1 or NaOH to obtain the desired pH under an atmosphere of Ar. The solutions were then degassed once more. An evacuated bulb of known volume was filled to the desired pressure of nitric oxide and then expanded into the evacuated bulb containing the ferrous chelate solution. Once the pressure had stabilized and was recorded (approximately 20 s), the solution was vigorously stirred. The pressure was monitored until it stabilized again, generally about 15 min. All pressure determinations were made with a Model 270 Baratron pressure gauge. After correcting for the water vapor pressure, the observed pressure drop was used to calculate the amount of nitric oxide absorbed by the solution. The solubility of nitric oxide in aqueous solutions (Armor, 1974) (1.95 X M/atm for pure water) was taken into account to determine the amount bound to the ferrous chelates. All work was done at 25 "C. Results and Discussion Reference experiments were done with nitric oxide expansions into water and sodium sulfate solutions to determine the accuracy of the method. These experiments yielded nitric oxide solubilitiesthat were in good agreement with values given by Armor (1974). Buffers were not used for pH control because of the possibility of interference with the ferrous chelate complex. Fe"(H,O), + NO. The ferrous ion concentration used ranged from 1.15 X to 5.0 x M, corresponding to ionic strengths of = 0.08-0.20. The solution pH was 3.3 in all runs. At a nitric oxide pressure of 1 atm, the hydrated ferrous ion did not completely bind with nitric oxide, due to the relatively small equilibrium constant for binding (Chang et al., 1982; Kustin et al., 1966). After correctingthe solubility of nitric oxide, the amount of nitric oxide absorbed was used to calculate the equilibrium constant, assuming 1:1Fe-to-NO stoichiometry. For K,, = [Fen(H20)5NO]/ [Fe"(HZO),][NO,], we obtained a value of 470 f 20 M-l, in excellent agreement with earlier results (Chang et al., 1982; Kustin et al., 1966). It does not appear that Fe11(H20)6 forms a dinitrosyl under these conditions, since the temperature jump experiments showed no evidence of a two-step decomposition (Chang et al., 1982; Kustin et al., 1966). Fe"(EDTA) + NO. All work was done with a ferrous M and an EDTA concenion concentration of 2.0 X 0 1987 American Chemical Society
Ind. Eng. Chem. Res., Vol. 26, No. 6, 1987 1233 1.0
I
*-,,Fe(E) (H20)6 \
Figure 1. Fraction of Fe(I1) bound to NO as a function of [IDA] to [Fe(II)] ratio a t pH 4.
tration slightly in excess of this value. The ionic strengths of the solutions were about p = 0.3. Experiments were done at pH 4.2, where the EDTAHZ2-form predominates, and a t pH 8.5, where the EDTAH3- form predominates. The equilibrium constant for the binding of nitric oxide to Fe"(EDTA) is large, about lo' M-' at 25 OC. With nitric oxide pressures of approximately 650 torr, the ferrous nitrosyl complex was found to have 1:l Fe-to-NO stoichiometry a t both pH conditions. I t was found that the ferrous nitrosyl complexes had very similar extinction coefficients a t 630 nm. This is consistent with work of Hishinuma et al. (1979) where it was reported that Fe"(EDTA) was a mononitrosyl and that the Fe-to-NO ratio was constant over pH 3 to 6. They reported the ratio dropped off below pH 3, presumably because of a decreasing fraction of EDTA in the EDTAH;- form. Ogura and Watanabe (1982), however, found 1:2 Fe-to-NO stoichiometry a t a low pH by using an indirect method. Fe"(NTA) NO. This complex also has a large reported equilibrium constant for the binding of nitric oxid, 2 x IO6 M-' at 25 "C (Lin et al., 1982). Ferrous ion and NTA concentrations used were about 0.020 M, with NTA slightly in excess of Fe(I1). Measurements were made at pH 7.5-8.5 and a t pH 2.5-3.5, with ionic strengths of p = 0.2-0.3. The complex was primarily Fer1(NTA3-)a t the high pH conditions and primarily Fe"(NTAH2-) at the low pH conditions. At the high pH conditions, 1:l Fe-to-NO stoichiometry was found for nitric oxide pressures from 150 to 800 torr. A t low pH, the Fe-to-NO stoichiometry was about 1:O.g. A t pH 3.1, the Fe-to-NO stoichiometry increased with increasing nitric oxide pressure over the range of 300-800 torr. With a final nitric oxide pressure of about 600 torr, the Fe-to-NO ratio decreased with decreasing pH, in agreement with Lin et al. The stoichiometry decrease with pH is presumably due to a decreasing fraction of NTA in the NTAH2- state. Ogura and Watanabe (1982) had found that FenNTA formed both monoand dinitrosyls in the pH 2-5 region. Fe"(1DA) + NO. IDA has a smaller stability constant for binding to Fe(I1) than NTA and EDTA (Ogura and Watanabe, 1982). It was necessary to have IDA concentrations in large excess to those of ferrous ion to ensure that the Fe(I1) was completely complexed to IDA a t pH 4. The predominant IDA species at pH 4 is IDAH-. Higher pHs were not studied due to the presence of the IDA2-species, which does not form a nitrosyl complex with ferrous ion (Griffiths and Chang, 1986). The IDA-to-Fe(I1) ratio was varied from 1:l to 20:1, and the Fe-to-NO stoichiometry approached 1:l a t the highest ratios, as shown in Figure 1. The final nitric oxide pressures were 630-640 torr. The ionic strength varied from p = 0.2 to 1.3, as no other salts were added to compensate for the variation in IDA concentration. By use of the values of the fraction of Fe(I1) bound to nitric oxide in Figure 1, the Keqfor Fe"(H,O), + NO, and the fraction of IDA as IDAH- at pH 4,a stability constant
+
3
4
I
I
Fe (II) (cit3-!/ ratio of Noabsorbed ,,'
5
6
---.
7
PH
Figure 2. Effect of pH on nitric oxide absorption by ferrous ioncitrate solutions and on ferrous ion-citrate complexes.
for Fe(I1) + IDA was determined, assuming that all of the Fe"(IDAH-) formed would bind with nitric oxide. A value of Kstab= [Fe"(IDAH-)]/[Fe"][IDAH-] = 1.0 M-' was obtained at pH 4. This does not agree well with the value of 6.3 X lo5 M-l given by Ogura and Watanabe (1982). It may be that the Fe"(IDAH-), complex alone absorbs nitric oxide, and the stability constant we have obtained is for the formation of this complex, whereas the previously reported stability constant is for Fe"(IDAH-), or for a different charge state of IDA. The assumption that almost all of the Fe"(1DA) would bind with nitric oxide appears to be valid, since the reported equilibrium constant for Fe"(1DA) binding nitric oxide is about 1.6 X lo4 M-' at pH 4 (Griffiths and Chang, 1986). Fe"(cit) NO. The Fe"(cit)(NO) complex was studied over a range of nitric oxide pressures and solution pH values with citrate-to-Fe values of 2:l and 1O:l. The ionic strength of the 2:l ratio solutions was p = 0.4-0.5 and that of the 1 O : l solutions was /I = 1.3. The ferrous ion conM. At the highest nitric oxide centration was 2.0-2.5 X pressures (700 torr) at pH 5 , the Fe-to-NO stoichiometry approached 1:l. At lower nitric oxide pressure, the amount of nitric oxide bound to Fe"(cit) displayed significant pH dependence. Hamm et al. (1954) reported that only citH2- and cit3- will bind with Fe(I1). By use of formation constants given by Hamm et al. (1954) for Ferr(citH2-)and Fe"(cit3-), the fraction of Fe(I1) present as each complex is plotted vs. pH in Figure 2. The ratio of nitric oxide absorbed to the total ferrous ion concentration is also plotted, after correction for the formation of Fe"(H,O),NO. The curve follows the curve of Fen(citH2-)between pH 3 and pH 4.5, but is slightly higher, suggesting that Fe(I1) + citH2- may form a complex with nitric oxide. The results were obtained with 0.025 M Fe(II), 0.050 M total citrate, and about 75-torr equilibrium pressure of nitric oxide. The decrease in the absorbed nitric oxide curve between pH 5 and 6 could indicate that Fe"(citH2-)N0 has a larger equilibrium constant for binding NO than Fe"(cit3-)N0. It could also be due to the formation of hydroxyl ion complexes. At moderate pH conditions, Fe"(cit3-) forms a complex with OH- that binds nitric oxide (Griffiths and Chang, 1986). In runs done at pH 8-9.5, the rate of nitric oxide absorption was much lower than in runs done a t pH < 7, and the ratio of absorbed nitic oxide t o total ferrous ion was also lower than runs done at lower pH. This could be due to the formation of a Fe"(cit3-)(OH), complex, where the second hydroxyl group blocks the nitric oxide from binding with the complex. This is supported by a titration curve of a solution of Fe(I1) and citric acid under a nitrogen atmosphere, shown in Figure 3. The dip in the curve between 3 and 4 equiv of base added agrees well with that observed by Hamm et al. (1954). A second dip occurs
+
Ind. Eng. Chem. Res. 1987,26, 1234-1238
1234
In conclusion, it appears that none of the ferrous nitrosyl chelates we studied exist as dinitrosyls under the conditions reported. The FeLIDA and F e k i t r a t e systems are complicated by the multiple charge states of the chelates and the formation of hydroxyl complexes in addition to nitrosyl complexes.
Fe(IL) + c i t r i c acid
Acknowledgment We appreciate the support and encouragement of Michael Persweig, Joseph Strakey, and John Williams. This work was supported by the Assistant Secretary for Fossil Energy, Office of Coal Research, Advanced Environment Control Division of the U.S. Department of Energy under Contract DE-AC03-76SF00098 through the Pittsburgh Energy Technology Center, Pittsburgh. Registry No. Fer1(H,0)5N0, 29966-86-7.
Literature Cited Armor, J. N. J. Chem. Eng. Data 1974, 19, 82-84. Chang, S. G.; Littlejohn, D.; Lin, N. H. ACS Symp. Ser. 1982,188, 127-152. 21
0
I
I
I
I
I
2
4 Equivalents of base odded
Figure 3. Titration curve of Fe(I1)
Griffiths, E. A.; Chang, S. G. Ind. Eng. Chem. Fundam. 1986,25,
J
356-359.
6
Ha",
R. E.; Shull, C. M.; Grant, D. M. J . Am. Chem. SOC.1954,
76, 2111-2114.
+ citric acid.
between 4 and 5 equiv of base, which could be attributed to the process
Hishinuma, U.; Kaji, R.; Akimoto, H.; Nakajima, F.; Mori, T.;Kamo, T.; Arikawa, Y.; Nozawa, S. Bull. Chem. SOC.J p n . 1979, 52, 2863-2865.
Kustin, K.; Taub, I. A.; Weinstock, E. Inorg. Chem. 1966, 5, 1079-1082.
Fe"(cit3-)(OH-) -tOH- = Fe"(cit3-)(OH-), Color changes also occured during the titration after 3 equiv of base was added. The solution turned from pale yellow to yellow-green and then to dark green as more base was added. This could also be due to the formation of the mono- and dihydroxy Fe"(cit3-) complexes. The color was different from that obtained in titrations of ferrous solutions without citrate, indicating that the Fen(OH), complex is responsible for it.
Lin, N. H.; Littlejohn, D.; Chang, S.-G. Ind. Eng. Chem. Process Des. Dev. 1982,21, 125-728. Littlejohn, D.; Chang, S. G . J . Phys. Chem. 1982, 86, 537-540. McDonald, C. C.; Phillips, W. D.; Mower, H. F. J. Am. Chem. SOC. 1965,87, 3319-3326.
Ogura, K.; Ishikawa, H. Electrochim. Acta 1983, 28, 167-170. Ogura, K.; Watanabe, M. Electrochim. Acta 1982,27, 111-114. Pearsall, K. A.; Bonner, F. T. Inorg. Chem. 1982, 21, 1978-1985. Received for review November 4, 1985 Accepted February 20, 1987
Application to Mixtures of the Peng-Robinson Equation of State with Fluid-Specific Parameters Zhong Xut and Stanley I. Sandler* Department of Chemical Engineering, University of Delaware, Newark, Delaware 19716
The vapor-liquid equilibria and volumetric behavior of mixtures have been calculated by using the Peng-Robinson equation of state with fluid-specific temperature-dependent parameters proposed earlier. Illustrative calculations show that this modification of the Peng-Robinson equation of state results in a great improvement in the simultaneous correlation of phase compositions and phase densities and excellent predictions of molar volumes in the one-phase region. The Peng-Robinson (1976) equation of state is a p = - -RT u - b U(U + b ) + b(u - b ) with a = rC,,R2T,2/P, and
b = +bRT,/Pc (1)
(2)
* Author
to whom correspondence should be addressed. Permanent address: Jiao Tong University, Sian, People's Republic of China.
0888-5885/87/2626-1234$01.50/0
(3)
In a previous paper (Xu and Sandler, 19871, we developed fluid-specific, temperature-dependent expressions for and +b from vapor pressures, saturated liquid volume, and supercritical volumetric data. The form of these expressions satisfies the classical critical conditions and, unlike the equations used by others (Yarborough, 1979; Morris and Turek, 1986) leads to +a and $b and their first derivatives, which are continuous in the critical region. Futher, the predictions for pure component molar volumes and vapor pressures obtained with these parameters were found 0 1987 American Chemical Society
