摘要:

J. CHEM. SOC. FARADAY TRANS., 1994, 90(4), 559-570 Speciation, Structural Characteristics and Proton Dynamics in the -Systems NH,NO, 1.5H20 and NH,NO, 1.5H20-(HNO, ,NH,F, NH,)-H20 at 5OoC Lars A. Bengtsson, Filip Frostemark and Bertil Holmberg* Inorganic Chemistry I,Chemical Center, University of Lund, P.O. Box 124,S-22100Lund, Sweden In order to obtain a good basis for exploring metal-ion complex formation in molten NH4NO3. 1.5H20, some fundamental characteristics of the pure hydrous melt and a number of compositions in the NH,NO,. 1.5H,0-(HN03, NH,F, NH,)-H,O system, have been investigated at 50"C.Several aspects have been taken into consideration, eg. thermodynamics of solvent autoprotolysis and HF formation, dynamics of proton exchange and structural properties.The acid dissociation constant of NH;, K, , and the equilibrium constant for formation of HF, K,,, were obtained from potentiometric measurements; K, = (2.2 & 0.2) x lo-' (mol kg-'), and K,, = 2160 & 40 (mot kg-I)-'. Results from "F NMR spectroscopy indicate that unprotonated fluoride, F-, probably exists as an H,NH+. * .F-ion pair in the solvent. The change in the "F chemical shift with increasing HNO, content in (NH,NO3-NH4F-HNO3). 1.5H20 verifies the conclusion from potentiometric data that HF is the only protonated fluoride species present. Raman spectroscopy and ' NMR experiments give clear evidence for an increased tendency to NH,+-..NO, ion-pair formation with decreasing water content in the systems NH,NO,-H,O.However, no loss of degeneracy of the internal v, and v, nitrate bands at 1380 and 718 cm-', respectively, was observed. The D,, symmetry of NO, seems to be preserved in the NH,N03. 1.5H20 melt. Results from Raman scattering, 'H NMR and l4N NMR experiments show significant changes in the spectra upon acidification with HNO, . These observations suggest an increase in hydrogen-bonding ability with increas- ing acidity. Results from large-angle X-ray scattering experiments on NH,NO,. 1.5H20 cannot be explained by a model comprising only interactions betwFen water molecules and ions. A residual contribution to the overall radial electron density distribution at 1.8 A is tentatively assigned to remarkably short N(NH,)-O(NO,) distances. 'H NMR spectroscopy shows a strong retardation of the proton exchange between NH; and H20 in the acidic region.The rate constant, k, , for the proton-exchange step H,N .HOH(OH,),-, + H,O -+ H,N . (OH,), + HOH, is estimated at (4.3 1.5) x lo7 s-'. This paper serves as a prologue to a number of investigations on metal fluoride and hydroxide complex formation in hydrous ammonium nitrate media at 50 "C.l Attention will be focused on the formation of polymetal complexes, i.e. com-plexes with formal composition [M, Xy+3, (m> 1, q a 1). Previous investigations in molten equimolar (K, Na)NO, at 240-280 "C provide evidence for the formation of dimetal complexes with a number of nontransition-metal ions, e.g. lead@) halide and hydroxide complexe~,~-~ alkaline-earth-metal fluoride complexes6 and lithium fluoride and hydroxide complexe~.~It has been emphasized that the formation of cationic species with high formal ionic charge should be enhanced by a purely ionic solvent, such as a molten ~alt.~,~ However, the high temperature normally required in work with molten anhydrous salts as solvents may cause crucial experimental problems.The introduction of a small amount of water to form hydrous melts allows lower working tem- peratures, which also enables the study of otherwise ther- mally unstable compounds, e.g. Bi(NO,), . For many so-called hydrous melts with less than 4-6 mol of water per mol of salt, the amount of water may be insufficient to form complete hydration shells around the constituent ions, and short-range Coulombic interactions like those in anhydrous melts prevail.lo Several investigations on metal-ion complex formation with chloride, bromide and iodide in hydrous nitrate melts, have been published over the years (see ref. 11-16 and liter- ature cited therein). Almost exclusively the main goal in these papers has been to explore the formation of anionic species, i.e. MXi-(n 2 1). In some cases the formation of hydrolysis products has been taken into account, at least qualitatively. To the knowledge of the present authors no investigation on complex formation with fluoride in such melts has been reported, probably due to experimental problems associated with hydrogen fluoride equilibria, as well as the more sophis- ticated data evaluation required.Ammonium nitrate holds a unique position in the proper- ties of salts in very concentrated aqueous solutions. The clas- sical works of Vollmar' and Narten' emphasize the striking similarity of the NH,f ion to the water molecule con- cerning mass, partial molar volume, bond angles and inter- atomic distances. NH,f and H20 form hydrogen bonds of about the same strength and the electrostatic forces exerted on their nearest neighbours are likely to be the same. Results from various spectroscopic and thermodynamic investiga- tions also suggest that the NH; ion fits well into the tetra- hedral local structure of liquid water.' Ammonium nitrate has been chosen as the solvent salt because of its very large solubility in water.The temperature of 50°C is appropriate in order to attain a hydrous melt with low water content, 1.5 H20 per NH,NO,, and it also allows the use of experimental techniques mainly based on com-mercially available equipment. The necessary prerequisites for our understanding of complex formation reactions between an ion such as fluoride and metal ions (with the inherent tendency to protolysis in hydrous media) obviously include a proper knowledge of the role of various protolytes of metal-free systems. This includes thermodynamics of solvent autoprotolysis and HF formation, structural charac- teristics and dynamics of proton exchange. These phenomena have been studied in the present work by use of several experimental techniques, such as potentiometry, multinuclear NMR, Raman spectroscopy and large-angle X-ray scattering.Experimental Chemicals Commercially available NH,NO, (Merck) and NH,F (Merck) used were of pA grade. Standard Karl-Fischer analyses showed the water content of NH,NO, to be 0.50 wt.%. The solvent melt, NH,NO,. 1.5H20,was prepared by mixing stoichiometric amounts of NH,NO, and doubly dis- tilled H20 at 50°C.NH,F was analysed according to the titration procedure by Lingane using La(NO,), (Merck, PA)." It was found to consist of 98.44 wt.% NH4F; the rest being H20. KF (Fluka, PA) was dried at 130°C for several days. NH,F, HNO, (65%, Merck, PA) and NH, (25%, Merck, PA) were used in preparing stock melts of various concentrations.NH, was analysed by acid-base titrations and was found to contain 22.5 wt.% NH, and water up to 100.0 wt.%. Fluoride-containing stock melts were stored in plastic containers in order to prevent the formation of fluoro- silicates as a result of hydrogen fluoride attack on glass. In particular, the NH, stock melts were handled with great care in order to avoid NH, losses. The reference melt used in the potentiometric measurements was prepared from AgNO, (Merck, PA) and stored in dark containers. All measurements were performed at 50°C. Concentrations are given in rnol kg-of NH,NO, unless otherwise stated. The melt composi- tion for all investigated systems was such that nH20/XnX = 1.5 (X = F-, NO,). EMF Measurements Apparatus The cell compartment consisted of a 100 ml Plexiglass vessel with thermostatted water encapsulating the inner part of the cell containing the test melt.The melt was stirred with a Teflon-coated magnetic stirring bar. Proton Fluoride Equilibria The change in fluoride-ion activity upon successive additions of (H,NH,)NO, * 1.5H20 to a NH,(F,NO,). 1.5H20 melt, was measured using a fluoride-ion-selective electrode with an LaF, membrane (Orion 94-09). The reference half-cell con- sisted of a Pyrex glass tube with a ceramic plug providing the contact between the internal and external melts. The internal compartment of the reference electrode was filled with an (Ag, NH,)NO,. 1.5H20 solution, 0.1 mol kg-' in Ag'. The cell may be described as NH,NO, * 1.5H20 AgNO, .1.5H20 II(0.1 mol kg-') F-innerNH4F * 1.5H2O (CF) membrane half cell The results of calibration series performed at various total concentrations of ammonia, are shown in Fig. 1. Fig. 1 clearly demonstrates the essential independence of ammonia concentration, indicating negligible formation of hydrogen fluoride species in the test runs. Hence, it holds that [F-] = CF for the calibration experiment. In the concentra- tion range 3 x lo-' < C,/mol kg-' < 0.4, the relation between the emf, EF ,and log[F-] is described by where EF = -588.9 0.2 mV and k, = 64.2 & 0.1 mV. The value of k, coincides well with the theoretical value at 50°C, [RT/F In 10],heof = 64.1 mV. After addition of a known amount of a stock melt of NH,(F,NO,).1.5H20 to the solvent melt, the cell was left for about 2 h in order to allow the temperature to stabilize. When stable emf readings of the cell were obtained, increasing amounts of a stock melt con- taining (H, NH,)NO, * 1.5H20 were added and the change in emf was recorded. A stable emf (within kO.1 mV) was obtained in <20 min after each addition. The total concen- trations used were lo-, < C,/mol kg-' < 5 x lo-, and < CJmol kg-' < 5 x J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 600 I I I I 550 c c\ 300 i350 0 1 2 3 4 5 -log(C,/mol kg-') Fig. 1 Results of calibration series for the fluoride ion-selective elec- trode at 50°C, performed at various total concentrations of ammonia; (0)CNH3= 0, (m) CNH3= 2.6 x mol kg-', (A) CNHJ= 9.2 x lo-' mol kg-' Evaluation of K, ,the Acid Dissociation Constant of NHf The changes in emf, upon addition of a stock melt containing NH, to the solvent test melt, were measured by means of a separate glass electrode (Metrohm 94-06) and an ammonia gas-permeable membrane electrode (Orion 95- 12).For the glass electrode the cell is schematically described as NH,NO, * 1.5H20 AgNO 1.5H20 /I(0.1 mol kg- l) In the concentration range lo-, < CH/mol kg- ' < 0.1, the emf of the cell follows Nernst's law. A typical test run is shown in Fig. 2. The relation between the emf, E, ,and C, is E, = Ei -kH log(CJmo1 kg-') (2) where Ei = 87.8 f 0.3 mV and k, = -65.9 & 0.2 mV. The observed Nernstian behaviour implies that the changes in the activity factors are negligible, and we may safely conclude that C, = [H'], representing the concentration of the hydrated proton H(H20),?.The ammonia electrode utilizes a hydrophobic gas-permeable membrane to separate the sample solution from the internal electrode solution. The internal filling solution I I I I I60 I E.-. LLII -60 > -**I J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 was replaced with a similar solution but more concentrated in NH4N03 (60.4 wt."/,), in order to eliminate the instability due to osmotic effects at the membrane. The cell is described as NH4N03* 1.5H20reference half-cell /I HNO, * 1.5H2O (CH) /INHZ ('NH3) ammonia gas- ammonia permeable membrane inner half-cell NMR Spectroscopy NMR measurements were performed at 50.0"C on a Varian Unity 300 MHz spectrometer, operating at 299.849 MHz for 'H, 21.665 MHz for 14N and 282.203 MHz for 19F.'H NMR spectroscopy was used to investigate the dynamic properties of test melts with varying hydrogen-ion concentration. The proton-exchange rates were studied in a number of melts with varying hydrogen ion and ammonia concentrations. Sample compositions are given in Table 1. The total concentrations given in Table 1 are based on weighed amounts of added HNO, and NH, stock melts, respectively 99.9% D,O (Sigma) contained in a glass capil- lary tube was used as external lock substance. 'H NMR chemical shifts are given relative to an external 0.06 g cmP3 aqueous solution of sodium 3-trimethylsilylpropanesulfonate, C,H, ,NaO,SSi.The proton signal of the reference sample was shifted 0.27 ppm to higher frequencies when the tem- perature was raised from 25 to 50°C. The lineshapes were analysed by the program DNMR520 and experimental rate 60 20 ,-20 E---.n = -60 LU' -1 00 -1 40 0.5 1.0 1.5 2.0 2.5 3.0 3.5 -log(CNH3/rnolkg-I) Fig. 3 Nernst plot of a calibration run at 50°C for the ammonia gas-permeable membrane electrode 56 1 Table 1 Total concentrations of HNO, and NH, used for the investigation of proton dynamics sample C,,,,/mol kg-' sample CNH3/mo1kg- 0.99 8 0 0.10 9 1.4 x 10-4 1.3 x lo-' 10 2.4 x 10-3 7.0 x 10-3 11 2.6 x 10-3 1.7 x 10-3 12 0.19 9.5 x 10-4 13 0.83 1.3 x 10-4 0 constants in the range 0.1 < kexp/s-' < lo5 were evaluated for the proton-exchange reaction by a visual fit of calculated spectra to experimental ones.14N NMR measurements were performed on melts with varying hydrogen-ion activity. The effects on I4N NMR spectra upon addition of F-were investigated. Measure- ments were also performed on aqueous NH4N0, solutions with varying water content. Neat CD,NO, was used as external reference. 19F NMR measurements were undertaken in support of the results obtained from the potentiometric investigation on the hydrogen fluoride complex formation. Samples with a constant total fluoride-ion concentration of mol kg- ' and HNO, concentrations ranging from lop3 to 0.16 mol kg -' were investigated.Measurements were also performed on samples with constant fluoride-ion concentration and increasing amounts of NH,NO,, ranging from dilute solu- tions to the hydrous melt composition, NH,N03 1.5H20. All samples were kept in 5 mm Plexiglass tubes, in order to avoid the formation of fluorosilicates. The resonance fre-quency of the 19F NMR spectra was measured relative to an external solution of neat CF,CO,H. Raman Spectroscopy The Raman spectra were recorded on a Bruker IFS 66-FRA 106 FT Raman spectrometer with a Ge diode detector and a low-power Nd :YAG laser (wavelength 1064 nm) providing the exciting radiation. The resolution was 4 cm-'. Sealed-off 5 mm standard NMR tubes were used as sample containers. A simple furnace consisting of an aluminium block electri- cally heated by internally mounted heating elements was employed.Large-angle X-ray Scattering, LAXS LAXS measurements on NH4N0, . 1.5H20 were performed with a GSD Seifert X-ray 8-8 diffractometer with a curved LiF monochromator of Johanson type, using Mo-Kcr radi- ation (A = 0.7107 A). The melt was kept in a Teflon con- tainer situated inside a closed glass vessel with 0.15 mm glass windows, in order to maintain a controlled atmosphere and to obtain a constant surface level. A slightly modified version of the previously described furnace and temperature control equipment was used.4 The monochromatized radiation scattered from the surface of the melt was measured in the range 0.65 < B/degrees < 58.3 in steps of s = 0.0335 (s = 47cA-l sin 0) by use of a scintillation detector.In the range 0.65 < 8/degrees < 3.2, lo4 counts were collected at each position. In the range 3.2 < B/degrees < 49.1, 4 x lo4 counts and in the range 49.1 < O/degrees < 58.3, 2 x lo4 counts were collected at each position. The observed inten- sities, lobs(8),were corrected for background radiation, Compton scattering, polarization and multiple scattering. After a normalization procedure the reduced intensities, i(s), J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 were obtained and Fourier transformed to give the reduced radial distribution function, D(r)-4nr2p0, where po is the average charge density of the atoms, All calculations were performed using the program KURVLR.2' The theoretical interaction parameters were refined using the least-squares program STEPLR.22 A more extensive description of the data treatment is given Results and Discussion Solvent Equilibrium For the NH,NO,.1.5H20 melt the following principal pro- tolytic equilibrium exists : NH: .(OH,), + (I + s -t)H,O=NH, *(OH,), + H(H,O),? Owing to the low overall water-to-salt-ratio the most likely values of I, s and t are 1 or 2. This reaction is characterized by the following equilibrium cons tan t (4) The ammonium ions and water molecules are present in large excess and their respective activities are therefore assumed to be constant.Hence, eqn. (4) may, to a first approximation, be simplified: K, = [NH,][H+] (5) The potentials of the ammonia and glass electrode were recorded after additions of NH, to the melt and [NH,] and [H'] were determined by use of eqn. (2) and (3). Based on a large amount of data the following value of the equilibrium constant was obtained: K, = (2.2 & 0.2) x (mol kg-')2 (6) Results of preliminary potentiometric measurements on metal-ion hydrolysis in slightly acidic solutions showed that the value of the product [NHJCH'] was equal to K, within the limits of error at each melt composition.' Proton Fluoride Equilibria Addition of a stock melt of (H,NH,)NO, -1.5H,O to a melt containing fluoride, causes a change in emf which may be attributed to formation of species such as HF, HF, or HzF+ according to H+ + F-eHF (11) H+ + 2F-e HF; (111) 2H' + F-e H2F+ (IV) etc. The formation reactions may be summarized xH+ + YF-=H,F;-~ (V) Assuming constant activity coefficients at constant water : anion ratio, the equilibrium constants corresponding to eqn.(V) may be expressed as (7) The total fluoride-ion concentration is given by CF = [F-] + [HF] + 2[HF,] + [H,F+] + Two separate cases can now be distinguished: EF(C, = 0) = E: -RT/F In CF (9) EF(CH)= 15: -RT/F ln[F-] (10) which gives AEF = EF(CH) -E,(C, = 0)= RT/F lIl(c~/[F-]) (1 1) By combining eqn. (7) and (8) we obtain C,/[F-] = 1 -k P1l[Hfl 28,,[H+][F-] + B21[H+I2 + ... (12) Clearly, a plot of CFICF-1 us.[H'] provides some quantita- tive information on the system: (i) Formation of only HF would result in a straight line with the slope /I1 ;(ii) forma- tion of polyfluoric species such as HF; would yield a depen- dence on CF; (iii) existence of HzF+ and other polyprotic species would be indicated by a curvature. A plot of C,/[F-] us. C, for three different values of CF is shown in Fig. 4. The fact that all data can be described by a single straight line clearly demonstrates that HF is the only species formed in significant amounts [case (i) above]. However, note that the approximation C, = [H'] is not a good one at lower values of C,. Such an approximation in the evaluation of the fluoride ion-selective electrode response might introduce errors in the interpretation of the data in that region.For the final computation of the equilibrium constant, the theoretical Nernst slope, total concentrations and emf data were used as input in the least-squares Fortran program EMFALL, which is based on a modified STEPIT s~broutine.~~The program minimizes the sum Z(AEexp -AECalJ2 in the search for the best model. The result of the computations is displayed in Table 2 along with some liter- ature data. Although K,,, the equilibrium constant for formation of HF according to eqn. (11), has been well characterized in dilute aqueous solution at relatively low ionic strength, no data seem to be accessible for ionic strengths larger than 8 mol dm-3 and for temperatures other than room tem-perat~re.~~.~~A smaller value of KHF is obtained in the present investigation as compared to results in concentrated NaClO, solutions (Table 2).24 This fact may be a result of the higher working temperature as well as the stronger solvent cation-fluoride interaction which destabilizes the hydrogen fluoride bond in the present solvent as compared to NaClO, solutions.As a result, a smaller value of KHF would be expected in the present study as compare to results of investi- gations in dilute aqueous NH4N03 solutions at room tem- perature. On the contrary, a larger value of KHF is obtained 300 I I I I I 1 I I A 250 -200 I kl2 150 100 50 0 0 0.02 0.04 0.06 0.08 0.10 0.12 0.14 C,/mol kg -' Fig. 4 C,/[F-] us. C, for different values of C,; (0)1.05 x(m) 3.20 x (A)4.68 x mol kg-' J.CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 Table 2 K,, in NH,NO, -1.5H20and some related systems ionic medium temperature/"C log K,," ref. ~ NH,NO,. l.SH,O 50 3.334 0.008 this work 0.5 rnol dm-3 NH,N03 25 2.84 28 0.5 rnol dm-3 NaCIO, 25 2.922 f0.003 26 2.0 mol dm-3 NaCIO, 25 3.108 f0.005 24 4.0 rnol dm-3 NaCIO, 25 3.539 f0.005 24 6.0 rnol dm-3NaC10, 25 4.058 0.006 24 8.0 rnol dm-j NaClO, 25 4.598 f0.010 24 PKHFexpressed in mol dm-3. (Table 2). This observation suggests that factors other than temperature and straightforward medium concentration effects are of importance for the stability of hydrogen fluoride complexes.This undoubtedly reflects the profound influence of the short-range ordering that becomes predominant as the molten salt characteristics begin to prevail over standard aqueous solution characteristics. The existence of other protonated species in aqueous solu- tion, mainly HF, and H2Ff, has been rep~rted.~~-~~ In an attempt to verify the apparent absence of such species in (NH,NO,-NH,F-HNO,) * 1.5H,O, "F NMR measure-ments were performed on melts with constant total fluoride- ion concentration, CF, and varying hydrogen-ion concentration. Only one single fluorine resonance signal was observed for these melts owing to the rapid exchange of fluo- ride between different sites. To a first assumption the single peak may be considered as emerging from a weighted average of the chemical shifts, bF-and oHF, characteristic of F- and HF, according to (13) a is the ratio [F-]/C,.The concentrations used and the observed "F chemical shifts, bobs, are displayed in Table 3. In order to evaluate the fractions of F- and HF, respec- tively, oF-and cHFwere estimated. The "F chemical shift of a strongly basic melt (NH, added in large excess) was taken as an estimate of c,-. A value of -16.0 ppm relative to CF,CO,H was obtained, which may be considered as the chemical shift of fluoride in an ammonium-ion environment. The other limiting value, bacid, was determined by extrapo- lation of bobs to C" ' = 0. This procedure is shown in Fig. 5. A further analysis of the 19F NMR data was made as follows.If HF is the only protonated fluoride species, it obvi- ously holds that baCid = bHF.Hence, [HF] and [F-] were estimated from eqn. (13), and the quotient QNMR = ([HF]/[F-])NMR was calculated for each value of C,. The free fluoride-ion activity for each sample was measured by Table 3 Total concentrations of HNO, and observed I9F chemical shifts CH/lO-mol kg- 156.2 -88.0 81.44 -87.7 51.17 -87.2 42.40 -87.0 31.18 -86.2 21.64 -84.6 10.43 -67.7 7.45 -54.2 4.86 -40.7 1.15 -22.0 --18.3 ~ ~~~ The fluoride ion concentration was within the range 8.6 d CF/10-, mol kg- < 10.0. ([HF]/[F-]),,, was formed. If the proper equilibrium model involves only the formation of HF, then the ratio, R, between the two quotients QpoJQNMR, should be equal to 1.The value of this ratio was found to be 1.0 & 0.2, which demonstrates the consistence between our analyses of potentiometric and "F NMR data in terms of a simple F--HF model. However, it should be emphasized that small systematic errors in oF- or UHF may cause substantial deviations from R = 1. Table 4 summarizes the results from the present work together with some pertinent literature data. The data displayed in Table 4 clearly demonstrate incon- sistencies in the chemical shifts reported, e.g. the 19F chemical shift of 'free', uncomplexed, fluoride. The word 'free' is some- what misleading since, of course, a 'free' or naked fluoride ion does not exist in a chemical environment owing to its high ba~icity.'~.~' However, the expression will be used throughout the present text to denote unprotonated fluoride.In order to investigate the influence of NH; ions, 19FNMR experiments were performed on samples with increasing con- centration of NH,NO, ranging from 0.1 mol dmP3 aqueous solution to the composition NH,NO,. 1.5H20.The fluoride- ion concentration, expressed as mol Fa (kg solution)- ' was kept constant. The results are shown in Fig. 6. From the diagram the 19F chemical shift of an F-ion in a dilute ammonium nitrate aqueous solution is estimated to be -42.0 ppm relative to CF,CO,H. This value seems to be reasonable compared with the literature data reported for dilute aqueous solutions (Table 4)., The shift of free F-moves towards higher frequency with increasing ionic strength which is in agreement with earlier reports, and may be explained either on the basis of an means of the fluoride ion-selective electrode, [HF] was calcu- increase in hydrogen bonding or on the formation of an ion lated as [HF] = C, -[F-1, and a similar quotient Q,, = pair, H,NHf- .-F-.32*36,37 Formation of protonated species, Table 4 Result of the 19FNMR measurements on the hydrogen fluoride system; literature data are summarized for comparison; chemical shifts are reported relative to neat CF,CO,H medium NH,NO,.1.5H20 aqueous solution aqueous solution aqueous solution aqueous solution aqueous solution Bu,NH:-,(F-/HF;) (n= 1-4y HC0,H OF -OHF ~HF~ ref.~ -16.0" -88.2 -this work -40.7' 33 -40.6' 34 -40.6' -83.2 -73.6 31 -40.6' -84.7 -73.1 35 -40.6' -61.1 -77.0 32 -44.6d --71.9, -79.6 86 -49.8' -83.6, -89.4 -49.8 -101.4 -67.4 51J a This value corresponds to the chemical shift of the fluoride ion probably existing as H,NH+-..F-. Fluoride present as F(H,O);. Molten liquids at room temperature except Bu,NHlHF; which is dissolved in CDCl,. n = 1. n = 4. 19FChemical shifts reported relative to C,F, in formic acid. The data were calculated using the expression a(CF,CO,H) = a(C,F,) + 84.4 ppm.'' 564 -85 t 1 -89l ' ' I I ' I ' 1 0 10 20 30 40 50 Ci'lkg mol-' Fig. 5 Extrapolation of oobs to C,' = 0 in (H,NH,)(F,NO,). 1.5H20 at 50°C 4035 n I I I I I I I 1 0 5 10 15 20 25 30 35 40 ' CNH4No3/mol (kg H,O) -Fig.6 Result of 19F NMR measurements on solutions with con- stant fluoride-ion concentration and increasing amounts of NH,NO,, going from a 0.1 mol dm-3 aqueous solution to NH,N03. 1.5H20. t = 50 "C. i.e. HF, HF;, would cause a shift towards lower frequencies which is in contradiction to what would be expected on the basis of simple first-order shielding arguments. Anion effects on the 19F chemical shift are, of course, much weaker than the cation effects, and an increase in nitrate concentration results in a contribution of <1% to the observed shift.36 Some authors assume that the effects in fluorine shielding are correlated to the strength of the hydrogen bond to the fluo- rine and ab initio calculations for the NH,-HF combination show a minimum-energy state for H,N..HF in which the proton of the hydrogen bond is residing closer to the fluorine than to the nitr~gen.~' In solid NH4F the hydro- gen bonding is between the ammonium cation and the fluo- ride anion, i.e. H,N-H..-F.40 The 3 ppm difference in chemical shift of fluoride in H,O and D20 solution indicates that hydrogen-bonding effects are of importance., 1*33*41 However, the interpretation of the results of "F NMR investigations on solvation of fluoride ions in mixed binary solvents rules out hydrogen bonding as the predominant factor controlling the shifts.42 This observation is consistent with the conclusions from an investigation on hydrogen fluo- ride in aprotic solvents, which suggest that other phenomena, such as a direct overlap interaction, affect the fluorine shield- ing as much as, or even more than, hydrogen bonding.43 In an early study of 19F chemical shifts in aqueous solutions containing F- and NH:, the decrease of fluorine shielding with increase of ammonium fluoride concentration was attributed primarily to an interaction with NH,f.37 An inter- J.CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 esting parallel can be drawn to the solvation of F-in for- mamide, HCONH, ,which represents a compromise between the formation of a hydrogen bond to the HCO-proton and direct N-F interaction with the nitrogen atom of the NH, group. The fluoride ion is located close to the amide A quantum-chemical study in the gas phase of the process NH, + HFeNHZF-shows that the interaction between NH, and HF is of a hydrogen-bonded nature.The system is most stable in the form of the neutral complex H,N. -.HF.45 However, in aqueous solution this reaction equilibrium is shifted towards the right owing to stabilization of the charged components by the medium. The presence of H20 molecules also facilitates proton transfer from HF to NH, It is .45746 therefore suggested that the ion pair, H3NH+...F-, is the most stable intermediate in an aqueous medium. Ab initio calculations carried out on hydrates of NH4F predict that ionic structures, e.g. H,NH+. --F-,are more stable than the neutral H,N. .HF, in the presence of solvating water.47 When more water is included into the calculations, the bridged4' and solvent-separates structures4' shown in Fig.7 become competitive in energy with the hydrogen-bonded ion pair. Rather surprisingly, results from relaxation time measure- ments on fluoride ions in aqueous solution, show that NH: ions do not have any different effect on the "F resonance, as compared to the alkali-metal ions. The dominant factor is a rejection of hydrogen bonding.48 It is not known which configuration of F- prevails in the present solvent melt, whether it is a hydrogen-bonded 'neutral' H,N.. .HF arrangement, an ionic form, H,NH+. .aF-, or a water-separated or bridged entity. It seems likely that a combination of these is the proper one. The neutral arrangement seems unlikely in the present hydrous melt owing to the strong ionic character of the medium. On the other hand, the large change in shift upon addition of NH,f ions to aqueous fluoride solutions is not consistent with a predominantly ionic NH: -F-interaction. The hydrogen-bonded ion pair, H,NH+...F-, or a bridged or water-separated structure are all likely candidates.The former seems most appropriate owing to the large change in 19F chemical shift observed when increasing amounts of NH,NO, are added to a melt containing fluoride. Such a model is consistent with the fact that both NH; and F-ions HI I H (b) Fig. 7 Structural models for the local fluoride environment in aqueous ammonium nitrate solutions; (a)ion pair, (b)bridged and (c) solvent-separated ion pairs J.CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 fit very well into the tetrahedral structure of the liquid water hydrogen-bonded net~ork.~~.~' The inconsistencies in oHFbetween the present medium and dilute aqueous solutions and formic acid may be due to the following facts. In one of the early reports glass tubes were used as container^.^^ Both fluoride and hydrogen fluo- ride solutions readily attack glass to produce SiFg-. The I9F chemical shift of the SiFg- ion occurs at higher frequency than those of HF and HF;. If a significant amount of SiFi-is present in the solutions, a lower value of the chemical shift of the observed average signal will result and errors appear in the evaluation of the chemical shifts of F-, HF and HF;. The chemical shifts of F-, HF and HF; in the investigation performed on formic acid are reported relative to C6F6 .51 This substance is not recommended as an internal standard owing to significant solvent shifts.52 A proper comparison with chemical shifts using CF3C0,H as reference is therefore not possible.Raman Spectroscopy Measurements Raman spectra of an unperturbed nitrate ion of D3h sym-metry ought to exhibit peaks from three fundamental vibra- tional modes: v,(A'J z 1050 cm-', v3(E')z 1380 cm-', v4(Ef)z 716 cm-'.53 The NHZ ion has symmetry, which gives rise to four Raman-active vibrational modes. The fre- quencies of these modes taken from solids are v,(A;) z 3040 cm-', v2(E) z 1680 cm-', ~3(T2) z 3145 cm-' and ~4(T2)x 1400 ~m-'.~~ In aqueous solution, or whenever water is present, 0-H stretching vibrations appear at about 3200- 3600 cm-'.53*55 Table 5 presents some literature data on Raman spectra of NH4N03 in aqueous solution, liquid ammonia solution, melt and the solid state. The interpretation of the Raman spectra may be difficult owing to band overlap. A tendency towards splitting of the asymmetric stretching vibration mode v3 and of the in-plane deformation v4 of the nitrate ion, may be ascribed to a lower- ing of symmetry from D3h to C,, or C,. Such effects are attributed to nitrate-water and nitrate-cation inter-action~.~~,~~,~~A more evident loss of degeneracy is the result of the formation of ion pairs, and monodentate and bidentate metal nitrate c~mplexes.~~,~~For the covalently bonded complexes also a low-frequency metal-oxygen stretching vibration mode Raman spectra of molten NH4N03 and ND4N03 indicate a splitting of the nitrate v3 fundamental mode on the basis of a curve-resolved fit of the broad asymmetric band centred at about 1370 cm-'.58 The two components at 1320 and 1410 565 cm-' were attributed to a splitting of the v3(N03) band. This result is in contradition to neutron diffraction data on molten ND4N03 .59 The interatomic distances found conclu- sively demonstrate that the NDZ ions conform to a tetra- hedral, and the NO, ions possess an equilateral triangular geometry. The NO, ions seem to retain their D,, symmetry in the melt, and consequently no loss of degeneracy of the E' mode of the NO, ion should be expected.Note that only one single band was reported in the 715 cm- ' region,s8 which normally is explained in terms of the absence of long-lived ion pairs. However, splitting of the v3 band was also reported in infrared studies on a 1 : 1 NH4N03 glass.60 That system shows a close resemblance to the ionic liquid. For the anhy- drous molten salt Av3 = 90 cm-' and for the glass Av3 = 82 cm-'. The v3 splitting of molten NH4N03 is interpreted as originating from a specific cation-anion interaction involving some hydrogen bonding.58 These results are in agreement with infrared spectra of solvated H30tN0, and NHZNO; in an argon matrix.60 The distorting ability of H30+ and NH,f equals that of the alkali-metal ions.The large Av3 values are suggested to arise from the effect of the diffusively charged H30+ and NHZ ions which are engaged in hydro-gen bonding with the NO; ion. The ammonium and hydro- xonium ions interact much more specifically with an NO, ion than do the alkali-metal ions, e.g. Li+. This may be an effect of hydrogen bonds being directional and specific, whereas Coulombic forces are not. Liquid ammonia, with a relative permittivity considerably lower than that for water, favours the formation of ion pairs. Some reports indicate contact ion-pairing between NH,f and NO, ions, based on loss of degeneracy of the v3 band Some authors report that the splitting of the v3 band is due to a hydrogen-bonding interaction with the solvent, and since the v4 bending mode is symmetrical in the NH,NO, solutions, contact ion-pairing is not e~ident.~~,~' The splitting of the degenerate v3 mode is ca.17 and 51 cm- less than the values reported in aqueous solution63 and melts, respectively.s8 This result indicates weaker hydrogen bonding in liquid NH, than in aqueous solution and NH,NO, melt. The absence of any measurable splitting of the NO, v3(Ef) or v4(E') modes has also been In this case no attempt was made to resolve the envelope because of overlap with the NHZ vk band. Aqueous ammonium nitrate solutions give little indication of disturbance of the anion symmetry by the cation and it is assumed that the ammonium ion, with a tetrahedral shape, fits well into the hydrogen-bonded network of water."A' Table 5 Observed Raman band positions (cm-') and assignments for some chosen NH,NO, systems reported in the literature 1052 1400 725 88 1048 1665 1420 720 3 100 1680 3220 1400 89 1042 1657 1336 707 2777 1700 3034 1436 61 1375 1040 1657 1325 710 2774 1705 3049 1379 64 1046 1332 712 62 1381 1044 1658 1320 71 5 3150 1677 58 1410 1043 1289 715 1418 90 1415 1461 1043 1655 1289 714 3140 1672 3192 1418 91 1414 715 3220 1445 3262 1462 1042 1655 715 3135 3223 1415 92 ~~ ~ Primed wavenumbers correspond to the NH,f ion. J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 0.10 0.08 0.06 0.04 0.02 0 3500 2700 1900 1100 300 ' wavenumber/cm -Fig. 8 Raman spectrum of NH,NO,.1.5H20at 50 "C Studies of relaxation processes in aqueous solution show that an increased concentration of NH,f has almost no effect on the vibrational relaxation time, zy, or on the rotational reori- entation time, 7or, for nitrate.66 The spectral properties of concentrated aqueous solution, where the immediate environ- ment of an NO; ion must include NH: as well as H20, are not appreciably different from those of dilute solution.55 Raman studies of vibrational dephasing processes suggest an inability of the NO; oscillator to sense the difference between NH:, NO; and H20 in its nearest-neighbour inter- action.67 The following sequence is found for the strength of this interaction: H,O z NH: z NO; < Na+ G Li+. Only at very high concentrations a small disturbance due to ammonium-nitrate interactions appears in the spectra in the form of band broadening.68 Fig.8 shows the Raman spectrum of NH4N0, 1.5H20 at 50°C. The two bands at 718 and 1048 cm-'correspond to the v4 and v1 internal nitrate vibrational modes, respectively. The band centred at about 1380 cm-' is assigned to v3(NO;) and vk(NHf). The appearance of a band at 1680 cm-' is due to the v; vibrational mode of the NHZ ion, but may also contain contributions from the first overtone of the Raman- forbidden fundamental vibration 2v2. The spectrum also shows the appearance of a broad asymmetric band at about 3140 cm-' which is attributed to both N-H and 0-H stretching vibrations.It has not been possible to resolve the components of the bands at 1380, 1680 and 3140 cm-'. The bandwidth at half height for the totally symmetric stretching vibration, v1 is 10 cm- ',which is about the same as reported in aqueous solution.69 However, a slight increase in band- width is observed for the v4 band at 718 cm- ' compared with the results in liquid NH, (17 and 13 cm-', respectively).61 The addition of NH4F and NH, to NH4N0,. 1.SH20 causes no significant change in the spectra. Addition of a 1 mol kg-' HNO, stock melt causes a small but significant 0.03 t 0.02 0.01 I I shift of the totally symmetric v1 vibrational mode of the NO, ion.The band in the 0-H and N-H stretching region is also influenced. These observations may be the result of an increase in hydrogen bond strength to the nitrate ion. For the sake of comparison, Raman spectra were also recorded for melts/solutions of composition NH4N0,. xH20 at 50°C, x = 3, 6 and 11. The spectra are shown in Fig. 9. The effects of increasing amounts of water in NH4N0,.xH20 are clearly displayed in the spectra. An increase of the anion-water and cation-water interactions with increasing water content is evident by the shift of the band in the 0-H and N-H stretching region, and also from the broadening and increased tendency towards split- ting of the internal v4 nitrate band. Considering the small amount of water present in NH4N03 * 1.5H20, this observa- tion implies a substantial increase in ion-ion interactions in the pure hydrous melt.Evidence for the formation of direct NHZ-NO; ion pairs does not exist since no splitting of the internal v4 nitrate band can be observed. This observation verifies the very small change in polarizability of the NO; ion exerted by the nearest-neighbour NH,f ions. The NO; appears to a larger degree to retain D,, symmetry in the hydrous melt than in dilute aqueous solution. Since small changes in the NO; ion symmetry are easily detected, the hydrous melt is a favourable solvent for detection of metal-nitrate interactions by Raman spectroscopy. LAXS Experiments The determination of the local structure of NH4N03 solu-tions by use of LAXS techniques is complicated by the physi- cal and structural resemblance between the ammonium ion and the water molecule, weak intermolecular interactions and also the small scattering power exerted by NH; and NO,.It has been suggested that the effect of ammonium salts on the water structure may be regarded as arising primarily from the anions alone,I8 and the first direct proof of NO;-H20 inter-actions in aqueous solution, provided by diffractometric investigations, was proposed by Caminiti et al. in a study of NH4N0, * 4.4H20.70A peak at 3.5 8, in the electron density correlation function was attributed to an N(N0,)-O(H20) distance. First-and second-order isotopic difference methods of neutron scattering have provided more accurate information about the local structure. A neutron scattering study on 12 and 18 mol kg-' aqueous solutions of ND,NO, demon-strates a direct interaction between the NDZ and NO; ions at the highest c~ncentration.~' An N(ND4)-O(N0,) distance of 2.15 8, is reported.Despite strong peaks due to intramole- cular atom-atom distances at 1.04 8, for the NDZ ion and 1.23 8, for the NO, ion, the rest of the total radial distribu- tion functions are relatively featureless as a result of the weak hydration. Neutron scattering studies of ND,NO, at higher 1 I I I I I I 3400 3200 3000 2800 775 750 725 700 675 650 waven umber/cm -' Fig. 9 Raman spectra of NH,NO,. xH,O at 50 "C,x = 1.5 (a),3 (b),6 (c), 1 1 (d) J. CHEM. SOC. FARADAY TRANS., 1994, VOL.90 concentrations, viz. 50 mol kg-', fail to resolve any N(ND,)-N(N0,) interaction below 3.6 This observation has also been made for the anhydrous molten salt.59 The results suggest a significant change in the structure of NH,NO, solutions in the concentration range 12 < CNH4N03/mo1kg-' < 50. This change in local structure is cor- related to the change in kinetics of the detonation process of ammonium nitrate at an H,O: NH,NO, ratio of 2: 1. In this respect, Adya and Neilson emphasize the need for a structural investigation of liquid NH,NO, .2H20.72To the knowledge of the present authors no such investigation has yet been performed. The nearest-neighbour cation-anion distance is reported to increase in the order NH,NO,(s), (3.82 < NH,NO,(l), (4.5 < NH,N03-D20(1), (6.1 A).', This effect upon melting NH,NO, is opposite to that of the alkali-metal nitrates, in which the cations move closer to the anion when the crystal melts.', Studies with the ND: cation, indicate that no extensive rearrangement seems to occur during the fusion process for molten ammonium nitrate, and hence the nitrate ions retain their D,, symmetry.59 The reduced radial distribution function for NH4NO3.1.5H20, is shown in Fig.lqa). The peaks at 1.18 and 2.14 8, are internal structure features of the nitrate ion. The peak at 3.17 8, and the shoulder at 3.70 8, are assigned to water-nitrate distances, O(H20)-O(NO,) and O(H,O)-N(NO,), respectively. The shoulder at 2.70 8, is assigned to O(H,O)-N(NH,) correlations.The broad band around 5 8, may contain contributions from NHZ-NO, con-tacts. The assignments are based on results of LAXS and neutron scattering measurements cited above. The structure model has been refined by means of a least- squares analysis and a comparison with experimental data is shown in Fig. lqb). The resulting difference curve clearly dis- plays deviations from experimental data at about 1.8 and 2.9 8,. This observation strongly indicates the inadequacy of a model based on only water-ion contacts and implies the exis- tence of direct ammonium-nitrate interactions. Oxygen atoms from water molecules are not likely to interfere since distances in this region are missing in more dilute solutions.70 The shoulder at about 1.8 8, in the reduced radial distribu- tion function may be assigned to O(N0,)-N(NH,) contacts.This distance is remarkably short compared with the value of 2.15 8, obtained in an 18 mol kg-' NH,NO, solution.71 The outcome of this comparison and also the fact that this peak is absent in a 50 mol kg-' aqueous solution of ammonium nitrate,72 definitely reveal the need for a critical assessment of the result of the present investigation. The distance of 1.8 8, is definitely too long for intramolecular interactions but seems on the other hand too short for intermolecular contacts. These facts inevitably suggest that the shoulder at 1.8 8, should be an artefact resulting from the data treatment or from unwanted scattering contributions to the overall inten- sity from the glass windows of the cell used.However, pre- liminary results from LAXS experiments in open containers on NH,NO,. 1.5H20 melts containing foreign ions still reveal the existence of a peak at about 1.8 A.' The data treat- ment was also slightly varied from time to time depending on the system investigated. Hence, we feel safe in concluding that the 1.8 8, peak is a physical reality and we suggest that it arises from a direct N(NH,)-O(N0,) interaction. An obvious preferred interaction geometry would be one with the oxygen atom of the nitrate triangle approaching the nitrogen atom at the centre of one of the triangular faces of the NH; tetrahedron, yielding a linear N(NH,)-O(N0,)-N(N0,) segment with tetrahedral H(NH,)-O(N0,)-H(NH,) angles.In such a configuration the N(NH,)-N(N0,) distance must be in the range 567 61 II I I I & -41Y 0, -61 -8 1 I I 0 2 4 6 8 10 r/A (b) i in QO h Fig. 10 (a) LAXS data: reduced radial distribution function for liquid NH,NO,. 1.5H20at 50°C. (b) D(r)-4nr2p, functions for NH,NO, . 1.5H20. Calculated peak shapes (-- - -), experimental curve(-) and the difference (---) are shown. d[N(NH,)-O(NO,)] + d[N-O(NO,)], i.e. (1.8 + 1.18) 8, = 2.98 A, which may contribute to the deviation at 2.9 8, in the difference curve shown in Fig. lqb). Unfortunately, owing to significant overlap of the peaks in the reduced radial distribu- tion function, no least-squares refinement yielding physically reasonable values of the temperature factor and the number of distances could be performed.It is not possible to derive an unambiguous structure model from LAXS data alone. A neutron scattering investigation of NH,NO, . 1SH20 might add valuable information. 14N NMR Measurements In the discussion, vide supra, on the structural properties of NH,NO, systems, it was noted that only small changes in ion-ion and ion-solvent interactions are observed for a hydrous melt compared with those of an aqueous solution and the anhydrous molten salt. Both the NHZ ion and the NO, ion retain an almost unperturbed tetrahedral and tri- angular geometry. 14N NMR experiments therefore are expected to show only small, if detectable, changes in chemi- cal shifts compared with results from aqueous solutions.Rota- tional correlation times for nitrate salts obtained from 14N NMR measurements indicate that the NHf ion does not associate with NOT even at high concentration^.^'*^^ The cited authors suggest that the weak complex-forming power of NH; is due to its tetrahedral shape, which enables it to fit too well into the hydrogen-bonded network without dis-rupting the water structure around the nitrate ion. This view was thoroughly discussed above, and is in agreement with results reported by others.' 7*1* Results of the present 14N NMR experiments on the pure solvent melt as well as melts containing NH,F, NH, and HNO, ,are shown in Table 6. A slightly more negative shift for the NO; ion is observed than for a 12.3 mol dmP3 saturated aqueous solution, corre- sponding to a molecular ratio of H20 : NH,NO, of 4.5 : l.77 This effect may be due to a combination of an increase in the NHi-NO; interaction, consistent with our Raman scat- tering results, and the higher temperature.Additions of fluoride of ammonia have no effect on the nitrate chemical shift. This is also in agreement with the results of the Raman spectroscopic studies. In the case of fluoride, a small shift of the ammonium 14N signal is observed, conceivably a result of the formation of an NH:-F-ion pair. Note that added NH4F amounts to only about one tenth of the total amount of NH4N03. A significant effect on the Raman spectra was observed upon additions of HNO, . This observation is also consistent with results of the present 14N NMR investigation. The chemical shift of the nitrate ion changes from -4.5 to -5.7 ppm. This trend is furthermore in agreement with results obtained in aqueous solution.77 Despite the strong tendency of H30+ to interact with water molecules, an increase in con- centration of HNO, and a decrease in water content, seem to produce NO;-H(H,O): ion pairs.The results may also reflect the increase in hydrogen bonding between nitrate and water. A similar interpretation may be valid for the results of the 'H NMR experiments on acidic melts in this work. Addi- tion of a 1 mol kg- ' HNO, stock melt causes a change in the shift of 0.55 ppm towards higher frequencies.The results of the 14N NMR experiments are thus in full agreement with the results of the Raman scattering measure- ments discussed above. Dynamics of Proton Exchange The kinetics of proton exchange of amines in water have been studied by Grunwald et u1.78-80 The mechanism of proton transfer may include either the formation of an amine hydrate according to H3NH+ . (OH,), + (r + s -t)H,O + H3N. HOH(OH,),- 1 + H(H,O),? ; k,, acid dissociation (VI) H3N. HOH(OH,),- + H20 + H3N .(OH,), + HOH; k,, proton exchange (VII) H3N * (OH,), + H(H,O),+ 4 H3NHt . (OH,), + (r + s + t)H,O; k-,, acid association (VIII) or for the formation of free amine H3NH+ + OH, + H,N + HOH; (1x1 HOH: + H,O + HOH: + HOH; fast (X) H3N + HOH: + H3NH+ + OH, (XI) J.CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 Owing to the observed dependence of exchange rate on [H '1 Grunwald et al. rejected the second mechanism and the proton transfer was suggested to occur according to the mechanism (V1)-(VII) including the formation of an amine- water complex, the so-called Swain-Grunwald mechanism." This suggestion has been supported by results of ab initio calculations suggesting that the process of proton transfer takes place in solution only if assisted by solvation or by a mechanism involving a third molecule.82 According to the proposed mechanism (V1)-(VIII), the experimental rate constant for proton exchange between NH: and water is given by The constants k,, kH and k-, refer to the three steps of acid dissociation, proton exchange and acid association described above.In the present investigation, variations of hydrogen-ion concentration produce changes in the proton-exchange rates, as monitored by 'H NMR experiments. In acidic melts where exchange rates are slow, a sharp singlet of the protons in the 'H,O site' and a 1 : 1 : 1 triplet of the NH: (14N, I = 1) ion appear. In the intermediate range, the protons move between these environments at a rate which is comparable to the frequency separation which causes successive broadening and merging of the signals. In alkaline melts where exchange rates are fast, only one signal appears. These observations are visualized in Fig. 11. For the present evaluation of rate constants of proton exchange in the rate region 0.1 d kexJs-' d lo5, computed spectra were compared with experimentally obtained ones.The remarkably slow exchange rate observed in an acidic environment is in agreement with earlier reports in aqueous sol~tion.'~-~~The observed dependence on [H'] for the experimental rate constant excludes the formation of free ammonia, and favours the ammonia hydrate model for the dissociation of NHZ ions in NH,NO,. 1.5H20. In the present systems, the acid dissociation constant of NH:, K, , is defined according to eqn. (5) in which K, may be expressed as K, = kJk-, according to eqn. (VI) and (VIII), which results in k, = [H+][NH,]k-, Table 6 Results of I4N NMR measurements obtained in the present study; some literature data are included for comparison; I4N chemical shifts are referred to neat nitromethane ~ medium bN03- 'JNH~+ ref.NH,NO, .ll H ,O -4.90 -360.39 this work NH,NO,. 6H,O -4.78 -360.39 this work NH,NO,. 3H20 -4.64 -360.45 NH,NO,. 1.5H,O -4.53 -360.59 this work NH4N03. 1.5H,0a -4.53 -. 36 1.1 9 this work NH,NO, * 1.5H,0b -4.53 -360.85 this work NH4N0,. 1.5H,O' -4.57 -358.25, -360.53, -362.94 this work (3 : 5 : 3) NH,NO,. 1.5H,0d -4.61 -355.67, -358.08, -360.53, -362.98, -365.44 this work (1 : 3 : 5 : 3 : 1) NH4N03. 1.5H20' -5.67 -355.63, -358.08, -360.53, -362.94, -365.39 this work (1:3:5:3:1) NH,NO, (adJ5.0 mol dm-3 NH,NO, -3.98 -4.64 -359.55 -358.96 77 77 in 2.0 rnol dm-, HNO, (aq) 0.5 < C,/mol kg-' < 1.9.CNH,< 0.95 mol kg-'. C, = 0.01 mol kg-'. C, = 0.1 mol kg-'. C, = 1.0 mol kg-'. 12.3 mol dm-j (satd, 30 "C). J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 1213 lo1'9 Fig. 11 Result of 'H NMR experiments on melts with various hydrogen-ion concentrations. The figures correspond to the concen- trations in Table 1. For the evaluation of the rate law for proton exchange, the rate of formation of the ammonia hydrate and its consump- tion according to eqn. (VII) and (VIII), may be used as a starting point; dCH,N. (OH2)Sl kexp = dt = kH[H3N* HOH(OH2),--k-aCH3N. (OH,)sICH(H,O),+I (16) By use of eqn. (15), this may be written in the equivalent form At lower values of [H+], when k,K,[H+] 9 k,, the relation between log kexp and -log([H+]/mol kg-') may be expressed as log kexp= log(kHK,) -log([H+]/mOl kg-') (18) The upper limit of [H'] for which the condition k, K,[H+] -' 9 k, holds, may be estimated from values of K, obtained in the present study, which is close to the value in aqueous solution, and k, and k, from investigations in dilute aqueous solutions. It is concluded that eqn.(18) may be safely used in the range [H'] < lo-, mol kg-'. A plot of log kexp us. -log([H+]/mol kg- ') for this concentration range results in a straight line of slope 0.96, which is near to the expected value of 1. The plot is shown in Fig. 12. The following value of k, was obtained: k, = (4.3 & 1.5) x lo7 s-'. This value of k, is almost four orders of magnitude smaller than values obtained for dilute aqueous solution.78 This observation may seem fairly reasonable owing to the hindrance of proton transfer by the large amount of nitrate ions present compared with dilute aqueous solution.A more complete investigation of the proton-exchange pro- cesses is out of the scope of the present paper and it should be emphasized that other effects such as temperature, vis- cosity changes and contributions to the observed rate from other proton-transfer reactions, may have a significant influ- ence on kexpand therefore also on k,. Concluding Remarks The solvent properties of NH,N03 1.5H20 have been explored and a good basis has been laid for the investigation of metal-ion complex formation in this medium.The fluoride ion-selective, ammonia and glass electrodes all show Nernst- ian behaviour in a wide concentration range. For systems based on NH,NO,. 1.5H20 the product [H+][NH,] is found to be constant, implying the existence 4.5 n 3.5 P)4 0)-2.5 1.5 3 4 5 -log([H+]/mol kg-') Fig. 12 Estimation of k,, the rate constant of proton exchange, log kcxp0s. -log([H+]/mol kg-') for lop6< [HC]/mol kg-I < of the following principal acid-base equilibrium NHZ *(OH,), + (r + s -t)H,OeNH, *(OH,), + H(H,O),f K, = [NH,][H+] = (2.2 & 0.2) x low9(mol kg-'), Potentiometric measurements demonstrate the formation of HF in the hydrogen fluoride system, K,, = [HF]/[F-][H+] = 2160 & 40(mol kg-')-'.19FNMR mea- surements confirm that HF is the only protonated fluoride species formed in significant amounts. Results from 19F NMR spectroscopy indicate that F- probably exists as an H,NH+...F-ion pair in the NH,(F,NO,). 1.5H20 melt. The triangular geometry of NO, is well preserved in NH,N03 1.5H20 as revealed by Raman scattering investi- gations. In acidic melts, the hydrogen-bonding ability is increased and H+-..NO, ion pairs seem to form. This con- clusion is in agreement with results from 'H and 14N NMR experiments. LAXS experiments reveal a shoulder at 1.8 A in the reduced radial distribution function which may be attributed to a direct N(NHl)-O(NO;) interaction. Results from 'H NMR experiments show a decrease in proton-exchange rate with increasing hydrogen-ion concen- tration.For the proton exchange H,N * HOH(OH,),-+ H,O + H,N. (OH,), + HOH the rate constant, k, was determined to be (4.3 f1.5) x lo7 s-'. Future work will be focused on the formation of polymetal complexes in the present medium. The molten salt-like properties of NH,NO, -1.5H20, which seem to be well pre- served, are of vital importance since the formation of highly charged polymetal complexes is known to be enhanced by the ionic medium. This work has been supported by a grant from the Swedish Natural Science Research Council. The Council is also grate- fully acknowledged for funding the Varian Unity NMR instrumentation. References 1 L. A. Bengtsson, F. Frostemark and B.Holmberg, to be published. 2 L. Bengtsson and B. Holmberg, J. Chem. Soc., Faraday Trans. I, 1989,85, 305. 3 L. Bengtsson and B. Holmberg, J. Chem. Soc., Faraday Trans. I, 1989,85, 317. 4 L. Bengtsson and B. Holmberg, J. Chem. Soc., Faraday Trans. I, 1989,85,2917. 570 J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 L. Bengtsson and B. Holmberg, J. Chem. SOC.,Faraday Trans. 1, 1990, 86,351. 51 C. Coulombeau, C. Beguin and C. Coulombeau, J. Fluorine Chem., 1977,9,483. 6 L. A. Bengtsson, F. Frostemark, B. Holmberg and S. Ulvenlund, 52 C. J. 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A. M. Walker, D. G. Lawrence, G. W. Neilson and J. Cooper, 26 G. T. Hefter, J. Solution Chem., 1984, 13,457. J. Chem. SOC.,Faraday Trans. I, 1989,85, 1365.27 G. T. Hefter, Polyhedron, 1984,3, 75. 72 A. K. Adya and G. W. Neilson, J. Chem. SOC.,Faraday Trans., 28 E. W. Baumann, J. Inorg. Nucl. Chem., 1969,31,3155. 1991,87,279. 29 K. Seppelt, Angew. Chem., Int. Ed. Engl., 1992,31,292. 73 R. W. G. Wyckoff, in Crystal Structures, Interscience, New York, C. Rieux, B. Langlois and R. Gallo, C. R. Acad. Sci. Paris, Ser. 2nd edn., 1964, vol. 2. 11, 1990,310, 25. 74 H. Ohno and K. Furukawa, J. Chem. SOC., Faraday Trans. 1, 31 K. Schaumburg and C. Deverell, J. Am. Chem. Soc., 1968, 90, 2495. 75 1978,74, 297. A. M. de P. Nicholas and R.E. Wasylishen, Can. J. Chem., 1987, 32 33 R. Haque and L. W. Reeves, J. Am. Chem. Soc., 1967,89,250. K. Radley and L. W. Reeves, J. Chem. Phys., 1971,54,4509. 76 65, 951.C. L. Perrin and R. K. Gipe, J. Am. Chem. SOC.,1986, 108, 1088. 34 C. Deverell, K. 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Jpn., 1981,54,1662. 46 A. N. Isaev and Yu. I. Khurgin, Izv. Akad. Nauk SSSR, Ser. Khim., 1991, 817. 91 H. C. Tang and B. H. Torrie, J. Phys. Chem. Solids, 1977, 38, 125. 47 P. Kollman and I. Kuntz, J. Am. Chem. SOC.,1976,98,6820. 92 D. M. Adams and S. K. Sharma, J. Chem. SOC.,Faraday Trans. 48 K. A. Alzewel. 2.Phys. Chem., Leipzig, 1974,255, 193. 2, 1981, 77, 1263. 49 A. K. Lyashchenko, G. V. Kokovina and A. S. Lileev, Zh. Strukt. Khim., 1987,28, 88. D. S. Terekhova, A. I. Ryss and I. V. Radchenko, Zh. Strukt. Khim., 1969, 10, 923. Paper 3/05193E; Received 31st August, 1993

ISSN:0956-5000    

DOI:10.1039/FT9949000559

出版商:RSC    

年代:1994

数据来源: RSC

摘要:

J. CHEM. SOC. FARADAY TRANS., 1994, 90(4), 571-574 57 1 Effect of Preferential Solvation on Reactivity of a Free Radical in Binary Solvent Mixtures Osamu Ito* and Hisanori Watanabe Institute for Chemical Reaction Science, Tohoku University, Katahira ,Aoba-ku, Sendai-980, Japan The absorption maxima of p-NH,C,H,S' and the rate constants and relative equilibrium constants for the reversible addition reaction of p-NH2C6H4S' with a-methylstyrene have been measured by a flash photolysis method in binary solvent mixtures with different compositions. In a polar-non-polar mixture, preferential solva- tion of p-NH,C6H,S' by the polar solvent was observed via the effect on the absorption maximum and the rate parameters. The linear free-energy relationship between the reaction rate constants and the relative equilibrium constants revealed that solvation of the reactant is more important than solvation of the transition state.Preferential solvation of a polar solute by the polar solvent in a polar-non-polar binary solvent mixture has been recog- nized from thermodynamic properties such as solubility and equilibrium constants. '-' Although it has been pointed out that the effect of solvation on the kinetic data is important for an understanding of the reaction process in solution, only a few studies on preferential solvation have been reported for ionic In the case of free-radical reactions, however, the effect of binary solvents on the reaction rates has not been reported, because of the non-polar character of the normal free redical, in addition to the difficulty in obtain- ing the rate parameters of the free-radical elemental reactions.In our previous studies,'-" it was found that flash pho- tolysis was an appropriate method for examining the effect of solvation on the free radical, because the reaction rate con- stant could be determined as an absolute value. In addition, the absorption maximum of the free radical could be observed at the same time by the same We found that the reactivity of the p-NH2C6H,S radical is quite sensitive to the solvent polarity, because of its dipolar ~haracter.','~ In the present study, the effect of binary solvent mixtures on the absorption maxima and reaction rate constants of p- NH,C,H,S' has been determined by flash photolysis in solvent mixtures with different compositions.In this reaction system, the linear free-energy relationship between the rate constants and the relative equilibrium constants, which could be obtained by analysing the reversible addition reaction, would be expected to afford information about the solvation of the transition state of the reaction. Experimental Materials (p-NH,C,H,S), was employed as a source of p-NH,C,H,S'. The disulfide was prepared from air oxidation of the corre- sponding thiol and was purified by recrystallization from ethanol. Commercially available a-methylstyrene (a-MSt) and all other solvents were used after purification. The non-polar solvents used were cyclohexane and carbon disulfide and the polar solvents were pyridine, ethanol and N-methyl pyrrolidin-2-one (NMP). Cyclohexane-pyridine and cyclohexane-NMP are expected to behave as standard polar-non-polar binary mixtures.The cyclohexane-ethanol mixture would be expected to show effects due to hydrogen bonding. CS,-pyridine and CS,-NMP systems were selected because some bituminous coals are quite soluble in the 1:1 (by volume) binary solvents at room temperature." (p-NH,C,H,S), cu. 5 x lo-' mol dmW3 was added to these binary solvent mixtures. This produced a sufficient con- centration Ofp-NH,C,H,S' (ca. 1 x lo-' mol drnp3) with one flash exposure to monitor the decay of the transient absorp- tion band in a cylindrical cell 10 cm long and 1 cm in diam- eter.' The oxygen concentration was controlled by the partial pressure.Apparatus The flash photolysis apparatus was of a standard design with two xenon flash lamps (xenon Corp. N-981C; with a half- duration of 10 ps). The photolysis light of wavelength 310-420 nm was selected by means of filters. The decay kinetics were followed by a photomultiplier detection system. Results and Discussion Absorption Maxima of p-NH,C,H,S' Fig. 1 shows the transient absorption spectra observed on flash photolysis of (p-NH,C,H,S), in cyclohexane, in pyri- dine and in their mixture. The transient absorption band observed in the visible region in each solvent system was attributed to p-NH,C,H,S', because the same absorption bands were observed on flash photolysis of p-NH2C6H,SH.9"3 The absorption peak shifts to longer wave- length with increase in the pyridine content of the mixture.This trend is in agreement with the general bathochromic shifts of the electronic transitions of polar solutes. The dipole moment in the excited state of p-NH,C,H,S' was determined 0.81 0.7 0.3 0.2 0.1 0.0 450 500 550 600 650 700 wavelength/nrn Fig. 1 Transient absorption spectra of p-NH2C,H,S'; (a) in cyclo- hexane, (b) in cyclohexane-pyridine (1 :1) and (c) in pyridine. The absorbance is shown immediately after flash of 1.0 x mol dm-3 @-NH,C,H,S), . 572 as 7.3 D, while that of the ground state was 4.3 D.13 A slight decrease in the absorption intensity was also observed; this is attributed not only to the hypsochromic effect but also to a decrease in the yield of the radical.For other binary solvent mixtures, similar tendencies in the absorption maximum and absorption intensity were observed. In Fig. 2 the variations of the absorption maxima of p-NH,C,H,S* are plotted against the mole fraction of the polar-non-polar binary solvent mixtures with cyclohexane as non-polar solvent. The line in each system shows the ideal case without specific solvation such as preferential solvation. For all binary solvents in this study, the absorption maxima move to a longer wavelength on addition of a small amount of polar solvent. This suggests that the polar solvent prefer- entially solvates p-NH,C,H,S' even though the content of polar solvent is less than that of the non-polar solvent.Fig. 2 shows that the extent of the preferential solvation is in the order NMP > ethanol > pyridine. The same tendency was observed when CS2 was used as non-polar solvent, although the absorption maximum in CS, was at a longer wavelength than that in cyclohexane. This suggests that CS, polarizes p-NH,C,H,S' more than cyclohexane; the polarizability also plays an important role. The extent of the preferential solva- tion in CS, on addition of a polar solvent was less than that in cyclohexane. Reaction Rate Constants Fig. 3 shows the first-order plots of the decay of p-NH2C,H,S' in the absence and in the presence of a-MSt in the binary solvent mixture of cyclohexane and pyridine.In the absence of a-MSt, p-NH,C6H4S' decays with second- order kinetics, thus, a deviation from the linearity of the first- order plot was seen [Fig. 3(a)]. This indicates that p-NH,C6H4S' decays predominantly by recombination to the disulfide. The low reactivity of p-NH,C,H,S* to oxygen was confirmed, since the decay did not change appreciably on the addition of oxygen to the ~olution.~ Since the arylthiyl radicals add reversibly to the alkene double bond, the decay rate of p-NH,C,H,S' was not accel- erated even in the presence of a-MSt in the degassed solution as shown in Fig. 3(b). In order to shift the fast equilibrium between the thiyl radical and the carbon-centred radical, it is necessary to add a radical trapping reagent which reacts selectively with the carbon-centred radical.Because oxygen has low reactivity to the thiyl radical and high reactivity to the carbon-centred radicals, oxygen may be appropriate for 18.5 1 1 18.1 r I 6 17.7 z ---2 I*E" 17.3 16.9 t I16.5 1 I 0.0 0.2 0.4 0.6 0.8 1.0 mole fraction Fig. 2 Spectral shift of the absorption maxima (imaX)of p-NH2C,H,S' in binary solvent mixtures of cyclohexane with (a)pyri-dine, (b) ethanol and (c) NMP J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 -2*o * -2.5 -3.0 :-3.5-this role. As shown in Fig. 3(c) and (4, the decay rate of p-NH2C,H4S' increases when oxygen and a-MSt are present in solution. The whole reaction scheme is represented by Scheme l.'*14 k' k'li +IOd 1k2 proxy radical Scheme 1 The decay rates of p-NH,C,H,S' depend on the concen- trations of a-MSt and oxygen.The rate constants in reaction Scheme 1 can be expressed as143's where [a-MSt] and kfirs,-ordcrrefer to the concentration of a-MSt and the first-order rate constant, respectively. The kfirst-ordervalue can be evaluated from the slope of Fig. 3(c) and (4.When the first-order plots are not linear, the contri- bution of the second-order decay due to the recombination of the thiyl radical can be eliminated by a graphic or computer simulation ' In Fig. 4, the plots of eqn. (1) are shown for cyclohexane- pyridine systems with five different ratios of pyridine to cyclo- hexane. The addition reaction rate constant (k,) can be evaluated from the intercept of Fig.4; the slope gives the value of Kk,[O,], where K = k,/k_, and [O2Isrefers to the oxygen concentration in the oxygen-saturated solution. The results for all the binary solvent systems studied are sum- marized in Table 1. J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 7.0I r( 1 6.01 n/"Id 5.0 X 4.0 ? In &c 3.0--.-6 2.0L 1 .o 0.0 0 1 2 3 4 5 6 1/[O,I, Fig. 4 Plot of eqn. (1) for the reaction of p-NH,C,H,S' with a-MSt in pyridine-cyclohexane binary solvent mixture; ratio = (a) 1 :0, (b) 3 :1, (c) 1 : 1, (d)1 :3 and (e)O:1 The log k, values are plotted against the mole fraction of each solvent for mixtures with cyclohexane, in Fig. 5. The values decrease with the increase in the mole fraction of polar solvent in each mixture.Preferential solvation by the polar Table 1 Rate parameters for Scheme 1 in binary solvent mixtures binary mixture f k'l mol-' dm3s-' (k-i/kz[Oz]s)/ mol-'dm3 Kk 2 ro2 I,/10-4mol-1 dm3 s-' cyclohexane- 0.00 23 0.14 170 pyridine 0.13 3.5 0.15 41 0.25 2.1 0.17 12 0.31 2.3 0.25 9.0 0.57 1.2 0.30 2.0 0.80 0.9 0.67 1.6 1 .00 0.8 0.85 0.9 cyclohexane- 0.00 23 0.14 170 ethanol 0.02 10 0.11 90 0.10 3.4 0.15 22 0.17 2.2 0.20 11 0.29 1.3 0.26 5.0 0.38 1 .o 0.22 4.5 0.64 0.7 0.21 3.2 0.85 0.7 0.20 3.3 1.00 0.8 0.26 3.5 cyclohexane- 0.00 23 0.14 170 NMP 0.01 4.2 0.18 24 0.03 1.7 0.25 6.9 0.05 0.71 0.17 4.2 0.11 0.45 0.26 1.7 0.22 0.38 0.32 1.2 0.27 0.35 0.38 0.93 0.53 0.25 0.40 0.63 0.77 0.22 0.44 0.50 1 .00 0.20 0.67 0.30 cs,- 0.00 9.2 0.46 20 pyridine 0.16 1.7 0.47 3.6 0.21 1.6 0.64 2.5 0.43 1.4 0.70 2.0 0.67 1.1 0.85 1.3 1 .00 0.80 0.89 0.9 CSZ-NM P 0.00 9.2 0.46 20 0.06 0.59 0.32 1.8 0.17 0.37 0.63 0.6 0.38 0.25 0.50 0.5 0.65 0.20 0.50 0.4 0.93 0.13 0.43 0.3 1.oo 0.10 0.50 0.2 x represents the mole fraction of the polar component, i.e.pyridine, ethanol or NMP. 573 5.0 ,-4.5 0 -4.0 3.5 3.0 0 0.2 0.4 0.6 0.8 1 mole fraction Fig. 5 Plot of log k, us. mole fraction of the binary solvent mixture of cyclohexane with (a)pyridine, (b)ethanol and (c) NMP solvent was observed in all solvent systems.In the case of the cyclohexane-NMP mixture, only 0.2 mole fraction of the polar solvent is needed to decrease the log k, value by 80-90%. The extent of preferential solvation by the polar solvent in cyclohexane increases in the same order as indicated by the absorption maxima in Fig. 2. In the addition reaction of p-NH,C,H,S' to a-MSt, it is assumed that the transition state has a dipole moment intermediate between the reactant and product which is the carbon-centred radical whose dipole moment may be negligibly small. This leads to a small reac- tion rate constant in polar media, because stabilization of the transition state by the polar solvent is less than that of the rea~tant.'O*'~~~ In Table 1, the k-,/k,[O,], values do not vary consider- ably on changing the concentration of the polar solvent. The k, values for the carbon-centred radicals with small dipole moment may not vary much on changing the solvent polarity.Thus, the variation of k-, between cyclohexane, pyridine, ethanol and NMP may be small. On putting k, = lo9 dm3 mol-' s-l and [O,],= lo-' mol dm-3, k-, can be calculated as ca. 3 x lo6 s-' which is large enough to estab- lish the fast equilibrium in Scheme 1 during a flash-lamp exposure of 10 ps duration. The small change in log k-,with solvent polarity can be understood from the less polar carbon-centred radical and the transition state of the reac- tion.On the other hand, the variation of Kk,[O,], with chang- ing solvent polarity is large. With an increase in the mole fraction of the polar solvent, the Kk,[O,], value tends to decrease. The decrease in log K implies an increase in the endo t hermici ty . In Fig. 6, the log k, values are plotted against log Kk,[O,], values to examine the linear free-energy relation- ship. A good linear relationship with slope of 0.75 can be seen; k,[O,],can be thought to be constant. A(1og k,) = aA(1og K) (2) In the equation above, o! = 0.75, which implies that the stabil- ization by solvation is one quarter of that of the reactant. In this study, we found a specific interaction within the binary mixture or with the solute. In the preferential solva- tion observed from the absorption maxima, a small amount of ethanol has a stronger effect than a small amount of pyri-dine in the mixture with cyclohexane (Fig.2), while a pro- nounced deviation cannot be found in log k, (Fig. 5). From Table 1 a specific interaction could not be deduced for CS,-pyridine and CS,-NMP. It was noticed that the rate 6.0 - 5.5 5.0 - c 0 4.5 - - 4.0 - - 3.5 - 3.0 2.5I 1 I I I I I I 2.5 3.0 3.5 4.0 4.5 5.0 5.5 6.0 6.5 ~og(~~,ro,l,) Fig. 6 Linear free-energy relationship (log k, vs. log Kk,[O,],) for the reaction of p-NH,C,H,S' with a-MSt in the binary solvent mix- tures shown in Table 1 constants and absorption maxima in the solvent mixtures with CS, tend to change gradually with time after mixing.Our data were determined immediately after mixing. Summary Preferential solvation by the polar solvent in a binary solvent mixture with a non-polar solvent was observed from its effect on the absorption maximum of p-NH,C6H,S' and its effect on the rate of the addition reaction with a-MSt. By compari- son of the reaction rate constants and absorption maxima, it is seen that preferential solvation of the reactant by polar solvent has a marked effect on the reactivity. From the linear free-energy relationship between the rate parameters evalu- ated from the reversible reaction, the preferential solvation of J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 the transition state was also evaluated to be one-quarter of that of the reactant p-NH,C,H,S'.The authors would like to express thanks to Prof. M. Matsuda and M. Iino and Ass. Prof. A. Watanabe of Tohoku University for useful discussions. References 1 Y. Marcus, Ion Solvation, Wiley, Chichester, 1985, ch. 7, pp. 185-217. 2 C. Reichardt, Solvents and Solvent Effect in Organic Chemistry, VCH, Weinheim, 1988, pp. 35-38. 3 P. Chatterjee and S. Bagchi, J. Phys. Chem., 1991,95,3311. 4 A. Chandra and B. Bagchi, J. Chem. Phys., 1991,44,8367. 5 K. Remerie and J. B. F. N. Engberts, J. Phys. Chem., 1983, 87, 5449. 6 V. S. Satri, R. W. Henwood, B. Behrendt and C. H. Langford, J. Am. Chem. SOC., 1972,94,753. 7 C. H. Langford and J. P. K. Tong, Acc. Chem. Res., 1977, 10, 258. 8 Y. Kondo, K. Uosaki, N. Tokura, Bull. Chem. SOC. Jpn., 1971, 44,2548. 9 0.Ito and M. Matsuda, J. Am. Chem. SOC.,1982,104,568. 10 0.Ito and M. Matsuda, J. Phys. Chem., 1984,88,1002. 11 0.Ito and M. Matsuda, Bull. Chem. SOC.Jpn., 1984,57, 1745. 12 H.Seki, 0.Ito and M. Iino, Fuel, 1989,68,837. 13 G. H. Morine and R. R. Kuntz, Chem. Phys. Lett., 1979,67,552. 14 0.Ito and M. Matsuda, J. Am. Chem. SOC., 1979,101,1815. 15 L. V. Natarajan, R. R. Lembke and R. R. Kuntz, J. Photochem., 1981, 15, 13. 16 E. F. Zwicker and L. I. Grossweiner, J. Phys. Chem., 1963, 67, 549. 17 G. L. Closs and B. E. Rabinow, J. Am. Chem. SOC., 1976, 98, 8190. 18 C. W. Fong, H. J. Kamlet and R. W. Taft, J. Org. Chem., 1983, 48,832. 19 0.Ito and M. Matsuda, Prog. Polym. Sci., 1992, 17, 827. Paper 3/04198K; Received 19th July, 1993

ISSN:0956-5000    

DOI:10.1039/FT9949000571

出版商:RSC    

年代:1994

数据来源: RSC

摘要:

J. CHEM. SOC. FARADAY TRANS., 1994, 90(4), 575-577 Limiting Partial Molar Volumes of Electrolytes in Dimethylformamide-Water Mixtures at 298.15 K Eugenio Garcia-Pafieda, Cayetano Yanes, Juan J. Calvente and Alfred0 Maestre* Departamento de Quimica Fisica , Facultad de Quimica , Universidad de Sevilla , 4 1012Sevilla, Spain Partial molar volumes at infinite dilution, V:, are reported for some 1 :1 electrolytes in dimethylformamide (DMF)-water mixtures covering the whole fraction range at 298.15 K. The results obtained show that the behav- iour of V; both for alkali-metal halides and for hydrophobic electrolytes is dependent on the added DMF. In the water-rich region, the alkali-metal halides exhibit small variations of VT, but two extrema, a minimum and a maximum, in the case of the hydrophobic electrolytes, are observed.In the DMF-rich region, a drastic decrease of V: for all electrolytes is exhibited. The results are discussed on the basis of ion-solvent and solvent-solvent inte ractions. The properties of aqueous solutions containing amides have solvent, respectively. A compilation of p, po and 4, are avail- received considerable attention since they have the peptide able as a supplementary publication.? linkage elements''2 and they are used as model compounds The application of the Redlich-Meyer equation24 was not to obtain information on biochemical systems. In relation to possible owing to the lack of values of the theoretical limiting aqueous mixtures of N,N-dimet hylformamide (DMF), several slopes, S,, in these mixtures.However, 4" were found to vary studies have been made in order to elucidate the mechanism linearly with m1/2over the concentration range investigated of the interaction of DMF with wziter (relative permit- (0-0.37 mol kg-'). The limiting partial molar volume of the cryoscopic and calorimetric mea~urements,~ ti~ities,~,~ ultra-electrolytes, VF =$:,was obtained by least-squares fitting sonic velocities,6 volumes and heat capacities,' NMR,8-'' viscosities",' ' and enthalpies of dilutionI2) where strong hydrogen bonding to the carbonyl group of DMF produces associates of the DMF-(H20), type. On the other hand, studies of electrolytes in aqueous mixtures of DMF (NMR,13 enthalpies of solution, transfer Gibbs energies' 7-19 and viscosities20.2I) have also been published, but limiting partial molar volumes of electrolytes in these mixtures have been restricted to tetraalkylammonium bromides.2 1.22 In this paper we report partial molar volumes at infinite dilution of 1 :1 electrolytes in aqueous mixtures of DMF covering the whole mole fraction range in order to obtain a better insight into ion-solvent and solvent-solvent inter-actions in mixed aqueous solvents.Experimental DMF (Merck, stated purity 99.8%, maximum content H20 0.05%) was dried over a thermally activated 4A molecular sieve prior to use. Ph,PC1 (Janssen Chimica, G.R., stated purity 99Y0) and NaBPh, (Merck, G.R., stated purity 299.5%) were dried for three days at 343 K in a vacuum desiccator.LiC1, NaCl, KC1, NaBr, KBr and KI were reagent grade (Merck) and were used after drying them overnight in an oven at 393 K. All salts were kept in a vacuum desiccator prior to use. DMF and water were degassed prior to making up solutions by weight. Measurements of densities were made using the apparatus and procedures described previ~usly.~~ Densities have an uncertainty (95% confidence limits) of +7 x g ~m-~. Results and Discussion Apparent molar volumes, 4,/cm3 mol-were calculated from solution densities using the standard expression M2 Po-P$,=-+-P WPOP where M, is the molecular weight of the electrolyte, m its molality; p and po represent the density of solution and of the results to the Masson equation: where S; is the experimental slope. Table 1 shows values of VF together with their 95% confidence limits (in parentheses).As can be observed, the experimental results obey the additivity rule within & 1.8 cm3 mol- '. In Fig. 1 we have plotted trends of VF with the mixed solvent composition for alkali-metal halides. At low DMF composition, up to xDMF= 0.3, the influence of added DMF does not produce significant changes on VF, except for LiCl where the effect of DMF begins to be noteworthy from xDMFz 0.17. On the other hand, as the DMF content increases, the effect of DMF produces a drastic decrease in VT up to pure DMF, taking even negative values for LiCl in the DMF-rich region. It is known that interactions between alkali-metal ions and dipolar aprotic solvents occur at the negative pole of the DMF dipole while the positive end of the dipole is sterically hindered from interacting with halide ions.In principle, this fact makes interaction of alkali-metal ions with both water and DMF possible, but that of halide ions is restricted to water only. However, on the basis of the Jones-Dole coeffi- cients of alkali-metal halides in water-amide mixtures, Woldan" has noted that in the water-rich region both alkali- metal ions and halide ions are solvated selectively by water, irrespective of the amide used. However, there is little evi- dence from standard functions of transfer for the preferential solvation of alkali-metal halides in dipolar aprotic aqueous solvents. On the other hand, Gallardo-Jimenez and Lille~~~ have studied enthalpies of interaction of alkali-metal halides with DMF in water in terms of the enthalpic virial second coefficient, hsaltPDMF,which is a measure of the water-mediated salt-DMF interactions.Note that in the case of our electrolytes, hsaltPDMFis negative both in the anionic and cationic series and becomes more negative as the ion size increases, except for LiCl, which seems to indicate that from a t Supplementary publication no. SUP 56985 (29 pp.), deposited with the British Library. Details are available from the Editorial Office. J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 Table 1 Partial molar volumes at infinite dilution, Vr/cm3 mol- ',for some electrolytes in xDMF-(1-x)water mixtures at 298.15 K X LiCl NaCl KCl NaBr KBr KI Ph,PCl NaPh,B 0.0 16.91" 16.62" 26.87" 23.48" 33.73" 45.6" 310b 276.4b 0.0196 0.0417 0.0955 0.1241 0.1 678 0.2227 0.2699 17.4 (0.4) 16.6 (0.4) 16.9 (0.4) 14.9 (0.4) 13.2 (0.3) 15.6 (0.3) 17.4 (0.3) 16.7 (0.3) 16.9 (0.9) 16.9 (0.5) 27.9 (0.3) 26.6 (0.1) 27.3 (0.1) 28.1 (0.3) 27.0 (0.4) 24.3 (0.1) 23.8 (0.3) 24.2 (0.3) 24.6 (0.5) 23.1 (0.7) 34.9 (0.7) 33.2 (0.3) 33.8 (0.3) 34.7 (0.2) 34.5 (0.6) 45.1(0.7) 46.4 (0.3) 47.2 (0.2) 48.1 (0.8) 47.5 (1.0) 310.1 (0.7) 309.5 (0.8) 307.3 (0.2) 306.9 (0.5) 305.9 (0.4) 307.9 (0.4) 307.7 (0.8) 274.5 (0.4) 278.5 (0.6) 288.9 (0.5) 296.9 (0.3) 299.4 (0.3) 0.3003 0.34 15 0.4232 0.4965 0.6598 0.6705 0.8299 1.o 8.6 (0.3) 5.4 (0.9) 1.7 (0.5) -4.4' 14.5 (0.4) 11.3 (1.0) C-5.9' 23.2 (0.4) 21.0 (0.9) 13.0' C- 21.5 (0.5) 17.0 (0.4) 14.0 (1.0) 6.6' 31.3 (0.4) 27.2 (1.0) 25.2 (1.2) 14.1' 44.1 (1.0) 41.7 (1.1) 35.0 (0.9) 30.5' 306.2 (0.3) 304.0 (1.O) 302.1 (0.6) 301.2 (0.2) 297.3 (1.1) 292.6 (0.8) 282.6 (1.0) 297.9 (1.0) 295.6 (0.4) 294.6 (0.6) 288.3 (0.3) 282.9 (1.0) 278.0 (0.5) 280' " Ref.24. Ref. 25. The solubility is too low to obtain accurate V?.'Ref. 26. Ref. 27. thermochemical point of view a favourable interaction obtained ionic partial molar volumes at infinite dilution, Vy, between the ions and DMF occurs. This is not in accordance of alkali-metal ions and halide ions in pure DMF from the with the idea that halide ions cannot interact with DMF ultrasonic vibration potentials method.Note that V? values because of steric hindrance at the positive pole of the DMF in water and in DMF are very similar for the halide ions, dipole. From the above considerations we could assume that except for the Br- ion, but those corresponding to alkali- alkali-metal ions and halide ions interact with DMF through metal ions differ markedly. This means that the difference their respective surrounding water cospheres in a water-rich observed in VT between water and DMF for alkali-metal region. The observed trends for VT could be attributed to halides in pure DMF is chiefly a consequence of cation- modifications in DMF-water interactions made by the pres- DMF interactions. ence of the ions in the mixed aqueous solvent. After this Fig. 2 shows trends of VT with mole fraction of DMF for water-rich region the water around the ions is progressively Ph,PCl, NaBPh, and Bu;NBr," the latter being included being replaced by DMF in the mixture and less water for comparison. In contrast to alkali-metal halides, the three remains to solvate the ions.This fact makes the interactions hydrophobic electrolytes exhibit significant changes of VT in predominantly of cation-DMF type, because halide ions are the water-rich region, with a minimum and a maximum sterically hindered from interacting with DMF molecules whose positions seem to be dependent on the type of electro- when the DMF content is high and, therefore, Vy behaviour lyte. Thus, the maximum is located at approximately the is mainly determined by the electrostriction effect of the same mole fraction of DMF for Ph,PCl and NaBPh,, alkali-metal ions.In this sense, Zana and Yeager3' have whereas for BuiNBr it extends up to the intermediate com- position region. From this extrema region the effect of DMF "I A '* 270 t 0.0 0.2 0.4 0.6 0.8 1.0 0.0 0.2 0.4 0.6 0.8 1.0XDMF Fig. 1 Variation of limiting partial molar volumes, V?, of electro- XDMF lytes with DMF composition for: KI (O),KBr (W), NaBr (A),KCl Fig. 2 Variation of limiting partial molar volumes, V?, with DMF (El), NaCl(A), LiCl(0) composition for: Ph,PCl (O),NaBPh, (A),BuaNBr (m) J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 effect with Ph,PCI and Bu2NBr electrolytes, and by the com- bined effects of the cation and the anion with NaBPh, elec-trolyte.I 550 -I I I I I produces a marked decrease of VF,similar to that for alkali- metal halides.The existence of a minimum and a maximum suggests that combined opposing effects between hydrophilic ion-solvent and hydrophobic ion-solvent interactions occur. Since the maximum is located at xDMFzz 0.27 for Ph,PCl and NaBPh, and considering that the Ph4P+ and BPh, ions have a similar size and comparable hydrophobic characteristics and, moreover, with negligible electrostriction effects due to their weak electric field, we could assign the behaviour of VT in the maximum as due to hydrophobic ion-solvent inter-actions. In contrast, the minima must predominantly be determined by the corresponding counter-ion-solvent inter-actions.In order to confirm that the assignments of both extrema are correct, we can cancel the respective contribu- tions of the C1- and Na' ions to VF of Ph,PCI and NaBPh, ,respectively, by considering the Ph,PBPh, electro-lyte, which can be achieved using the additivity rule in Table 1 according to the expression V?(Ph,PBPh,) = V,"(Ph,PCl) + V,"(NaBPh,) -V,"(NaCI) (3) In Fig. 3 we have plotted the dependence of VT on the mole fraction of DMF for Ph,PBPh,. The plot shows that the minimum has disappeared and only the maximum at xDMF z 0.27 prevails. On the other hand, the drastic decrease of Vr with increasing content of DMF for the three hydrophobic electrolytes (Fig. 2) has been attributed in the case of Bu:NBr22 to structural interactions between this and the DMF.Nevertheless, on the basis of the above considerations for alkali-metal halides, the behaviour of Vr in the DMF-rich region must be determined by the hydrophobic cation We wish to thank DGICYT of Spain for financial support (Proyecto no. PB91-0605). References 1 G. Somsen, Pure Appl. Chem., 1991.63, 1687. 2 G. R. Hedwig, T. H. Lilley and H. Linsdell, J. Chem. SOC., Faraday Trans., 1991,87, 2987, and references therein. 3 G. Douheret and M. Morenas, C. R. Acad. Sci., Ser. C, 1967, 2, 729. 4 R. Reynaud, C. R. Acad. Sci., Ser. C, 1968,266,489. 5 J. Bougard and R. Jadot, J. Chem. Thermodyn., 1975,7, 185. 6 F. Kawaizumi, M. Ohno and Y. Miyahara, Bull. Chem. SOC. Jpn., 1977, 50,2229. 7 C.De Visser, G. Perron and J. E. Desnoyers, J. Chem. Eng. Data, 1977, 22, 74. 8 Y. I. Mitchenko, V. A. Fenin and E. P. Krasnov, Russ. J. Chem., 1978,52, 163. 9 V. Zelano, 2.Phys. Chem. N. F., 1983,138,31. 10 C. M. Kinart, W. J. Kinart and L. Shulski, Pol. J. Chem., 1985, 59, 597. 11 C. Della Volpe, G. Guarino, R. Sartorio and V. Vitagliano, J. Chem. Eng. Data, 1986, 31, 37. 12 R. H. Wood and L. H. Hiltzik, J. Solution Chem., 1980,9,45. 13 A. Fratiello, R. E. Lee, D. P. Miller and V. M. Nishida, Mol. Phys., 1967, 13, 349. 14 S. Taniewska-Osinska and A. Piekarska, Bull. Acad. Pol. Sci., Sir. Sci. Chim., 1978, 26, 613. 15 L. Thakur and R. Prasad, Indian J. Chem., 1980, 19A, 520. 16 W. J. M. Heusvelsland, C. de Visser and G. Somsen, J.Chem. SOC., Faraday Trans. I, 1981,77, 1191. 17 S. P. Rudra, B. P. Chakravarty, K. Kundu and I. N. Basu-Mallick, Z. Phys. Chem. N. F., 1986,150,211. 18 E. A. Gomaa, Thermochim. Acta, 1989, 142, 19. 19 K. Das, K. Bose and K. Kundu, Electrochim. Acta, 1981,26,479. 20 B. N. Prasad, N. P. Singh and M. M. Singh, Indian J. Chem., 1976,14A, 332. 21 N. C. Dey, G. Kumar, B. K. Saikia and I. Haque, J. Solution Chem., 1985,14,49. 22 W. J. M. Heuvelsland and G. Somsen, J. Chem. Thermodyn., 1977,9, 231. 23 C. Yanes, P. Perez-Tejeda, E. Garcia-Paiieda and A. Maestre, J. Chem. SOC., Faraday Trans., 1992,88,223. 24 F. J. Millero, in Water and Aqueous Solutions. Structure, Ther- modynamics and Transport Processes, ed. E. A. Home, Wiley, New York, 1972, ch. 13. 25 C. Jolicoeur, P. R. Philip, G. Perron, P. A. Leduc and J. E. Des-noyers, Can. J. Chem., 1972,50,3167. 26 F. Kawaizumi and R. Zana, J. Phys. Chem., 1974,78, 1099. 27 M. R. J. Dack, K. J. Bird and A. J. Parker, Aust. J. Chem., 1975, 28, 955. 28 M. Woldan, Z. Phys. Chem. N. F., 1986, 150,201. 29 M. A. Gallardo-Jimenez and T. H. Lilley, J. Chem. SOC., Faraday Trans. I, 1989,85,2909. 30 R. Zana and E. B. Yeager, in Modern Aspects of Electrochem-istry, ed. J. O'M. Bockris, B. E. Conway and R. E. White, Plenum Press, New York, 1982, vol. 14, ch. 1. Paper 3/02039H; Received 8th April, 1993

ISSN:0956-5000    

DOI:10.1039/FT9949000575

出版商:RSC    

年代:1994

数据来源: RSC

摘要:

J. CHEM. SOC. FARADAY TRANS., 1994, 90(4), 579-582 579 Ultrasonic Velocities and lsentropic Compressibilities of some Tetraalkylammonium and Copper(1) Salts in Acetonitrile and Benzonitrile Jasbir Singh, Taranjit Kaur, Vazid Ali and Dip Singh Gill* Department of Chemistry, Panjab University, Chandigarh- 160014,India Ultrasonic velocities and densities of Bu,NBPh,, Pr,NBPh,, Et,NBPh,, Bu,NCIO,, Pr,NCIO,, Et4NCI0,, Bu,NI, CuCIO, .4CH,CN and [Cu(DMPhen),]CIO, in acetonitrile and benzonitile have been measured at 298, 308 and 318 K. The isentropic compressibilities (K~)and apparent molal compressibilities (K~)of various electrolytes have been evaluated. The limiting apparent molal compressibilities (K:) for all electrolytes have been obtained by extrapolation of the plots of IC+vs.square root of molality (m’I2)and split into limiting ionic compressibilities (I&). IC:* is positive in benzonitrile except for Et,N+ and is negative in acetonitrile except for Bu,N+, Ph,B-and Pr,N+. The results indicate stronger solvation of ions in acetonitrile than in benzonitrile. Compressibility data are usually needed for predicting the Results and Discussion pressure dependence of equilibrium properties of electrolyte solutions. Precise compressibility data of electrolytes in non- The ultrasonic velocities (u) as a function of the molality of aqueous solvents are rare1*2 and for copper(r) salts they are the electrolytes in AN and BN are shown in Fig. 1 and Fig. 2. completely lacking. In order to investigate the compressibility The ultrasonic velocities vary linearly with molality.In AN behaviour of some tetraalkylammonium and copper(1) salts in acetonitrile (AN) and benzonitrile (BN), we have undertaken ultrasonic velocity and density measurements of solutions of tetrabutylammonium tetraphenylborate (Bu,NBPh,), tetra-I propylammonium tetraphenylborate (Pr,NBPh,), tetraethyl-1296 ammonium tetraphenylborate (Et,NBPh,), tetrabutyl-ammonium perchlorate (Bu,NClO,), tetrapropylammonium 1292 perchlorate (Pr,NClO,), tetraethylammonium perchlorate (Et,NClO,), tetrabutylammonium iodide (Bu,NI), copper(r) 1288 perchlorate tetraacetonitrile (CuCIO, .4CH3CN) and bis( 2,9-dimet h yl-1,lO-phenant hroline)copper(i) perchlorate 1284 ([Cu(DMPhen),]ClO,) at 298,308 and 318 K. ~~1280 0 0.05 0.10 0.15 0.20 0.25 0.30 Experimental 1259 Ultrasonic velocity measurements were carried out at 2 MHz r 1254 Iwith an ultrasonic time intervalometer (model UTI-101) from fn Innovative Instruments, Hyderabad using a pulse-echo E 1250 \ 5overlap technique.Density measurements were made using a 1246precision densimeter (Anton Paar model DMA-60 with exter- nal measuring cell-602). The absolute accuracy of the sound 1242velocity measurements as reported before3 was better than 2 parts in lo4 and for density measurements it was & 1 x lo-’ 1238g cm-3. The ultrasonic velocities in pure AN at 298, 308 and 0 0.05 0.10 0.15 0.20 0.25 0.30 r 1318 K were 1280.8, 1239.6 and 1200.4 m s-l and in BN were 1418.0, 1382.5 and 1349.1 m s-’.The densities of AN at the 1219 corresponding temperatures were 0.776 851, 0.765 813 and 0.752717 g cmP3 and of BN were 1.00034, 0.991 683 and 1215 0.982945 g cm-3. A comparison of our all values with others 1211cannot be made, but the sound velocity in pure AN and its density at 298 K are in good agreement with the literature 1207 values. q4 AN (99% pure, E. Merck, India) and BN (Puriss, >99% 1203 GC, Fluka Chemika) were purified as reported The sources, grades and methods of preparation/purification of all 1199 1 0 0.05 0.10 0.15 0.20 0.25 0.30 the electrolytes used are given previously.6-8 A range of con- m/mol kg-’ centrations of the electrolytes in AN and BN was produced Fig. 1 Ultrasonic velocity us.molality in AN at (a) 298 K, (b)308 Kby diluting stock solutions of appropriate concentrations. In and (c) 318 K. *, Bu,NBPh,; 0, Pr,NBPh,; 0,Et,NBPh,; x,all cases the measurements were repeated to obtain repro- Bu4NC10,; V, Pr,NClO,; X, Et,NClO,; Q, Bu,NI; 0,ducible results. CuClO, .4CH ,CN ;A, Cu(DMPhen),ClO, . 580 1440 1436 1432 1428 1424 1420 1416 0 0.1 0.2 0.3 0.4 1380 I II 0 0.1 0.2 0.3 0.4 rnlmol kg-Fig. 2 Ultrasonic velocity us. molality in BN at (a) 298 K, (b)308 K and (c)3 18 K. Symbols as in Fig. 1. the ultrasonic velocities are smaller than those in BN. The sound velocities of tetraalkylammonium tetraphenylborates are larger than those of the corresponding perchlorates and Bu,NI in both solvents.The ultrasonic velocity of CuC10,.4CH3CN is smaller and that of [Cu(DMPhen),]ClO, almost comparable to those of tetra- alkylammonium salts in AN and BN. The ultrasonic velo- cities decrease with increasing temperature. The isentropic compressibility of each electrolyte was cal- culated from the ultrasonic velocity and density (p) using the equation : KS = 1/(U2P) (1) Fig. 3 and 4 show K~ as a function of molality. In contrast to the ultrasonic velocities, the isentropic compressibilities decrease non-linearly with increasing salt concentration and increase with temperature. The apparent molal volumes, V, and apparent molal isen- tropic compressibilities, K+ of all the electrolytes at 298 K have been calculated using the equations and (3) J.CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 79 1 I 78 77 76 75 1 I J 0 0.05 0.10 0.15 0.20 0.25 0.30 85 r I Z 83n 0I 0381 79 0 0.05 0.10 0.15 0.20 0.25 0.30 93 1 92 91 90 89 80 87 86 0 0.05 0.10 0.15 0.20 0.25 0.30 rn/mol kg-Fig. 3 Isentropic compressibility us. molality in AN at (a) 298 K, (b) 308 K and (c) 318 K. Symbols as in Fig. 1. where m is the molality, M is the molar mass (g mol-I) of the solute, K,(sin) and xS(1) are the isentropic compressibilities of the solution and pure solvent and &in) and p(1) are the densities (g cm-3) of the solution and the pure solvent. The limiting apparent molal compressibilities (@ of the electrolytes were obtained from the extrapolation of the linear plots of K, us.m112 by the least-squares method using the equation: K, = K$ +-A,m'I2 (4) Plots of K, vs. m112are shown in Fig. 5. The extrapolated IC:and A, values of eqn. (4) are reported in Table 1. The maximum uncertainty in the IC~values is f3 x lo-, cm3 mol -bar -I. Table 1 shows that K$ values in AN are negative for Et,NBPh,, Pr,NClO, , Et,NCIO, , CuClO, -4CH3CN and [Cu(DMPhen),]ClO, while for the other electrolytes they are positive. In BN K: values are positive for all the solutes. Our K: value for Et,NClO,, -91 x lo-, cm3 mol-' bar-', in AN is in good agreement with the value, -92.4 x lo-, an3 mol-' bar-', reported by Davidson et al.' Table 2 lists the pairwise differences of K$ values among a series of tetraalkylammonium salts to verify the additivity rule.The difference between appropriate pairs of electrolytes does not exceed 3 x lo-, m3mol-' bar-', which is within the limit of our experimental error. This verifies the additivity rule of compressibilities (K: values) in AN and BN. J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 49.8 I 1 0 0.1 0.2 0.3 0-4 r 1 52.8 52.4 52.0 I b 51.8 n 51.2 0 7 50.8 '50.4 1 50.0 1 Id 0 0.1 0.2 0.3 0.4 55.8 55.4 55.0 54.6 54.2 53.8 -53.4 There is no reliable and direct method of breaking K: values into ionic components. Some approaches have been The method' already used in AN is based upon K~(P~,B-) 0. This approach is not appropriate as the = Ph,B-ion is large (0.535 nm)," even larger than the Bu,N+ ion (0.50 nm)I3 and therefore its compressibility contribution cannot be taken as zero.Millero" has split the partial molar volumes of electrolytes into ionic components using Ph,AsBPh, as a reference electrolyte on the basis of v:Ph4 As + )/'gh4 B -) = $Ph4 As + )/':Ph4 B -) A similar model based on Bu,NBPh, as a reference electro- 58 1 150 100 50 0 -50 --100 n 7--150 I I I 1 I 1E I 0.1 0.2 0.3 0.4 0.5 0.6 m *g 200 I2--. Yi 100 0 -1 00 I I I I Fig. 5 Apparent molal isentropic compressibility vs. square root of molality at 298 K in (a) acetonitrile and (b) benzonitrile. Symbols as in Fig. I. lyte for splitting the viscosity B coefficients of electrolytes into the contributions from individual ions was suggested by Gill and Sharma.', This model has been used in the present work.The K:* values obtained in this way are listed in Table 3. The most striking features of the ionic K:* values in AN are the large positive values for Bu,N+ and Ph,B- ions, the very small value for Pr,N+ and the negative values for I-, ClO,, Et,N+, Cu+ and [Cu(DMPhen),]+ ions. In benzoni- trile, the ionic K:* values for all ions except Et,N+ are posi- tive. They are relatively large for Bu,N+, Ph,B-and [Cu(DMPhen),]+ and relatively small for Pr4N+, ClO, and Cu'. The large rc:* values arise either from some free space present in the solution between the ions and the solvent mol- ecules or from conformational changes in the structure of the large R,N+ cations. For a large R4N+ ion such as Bu,N+, the large K:* value can be explained due to changes of con- figuration from long chains into a coiled or spherical configu- ration by the application of pressure from the ultrasonic wave. This effect should decrease with decreasing size of the R,N+ cation.This is seen from the results of Table 3 as ionic Table 1 Limiting apparent molal isentropic compressibilities and the slope (A,) of eqn. (4) for some tetraalkylammonium and copper(1) salts in AN and BN at 298 K AN BN electrolyte K$/~O-, cm3 mol-' bar-' 104~~ K$/~O-, cm3 mol-' bar-' 104~, Bu,NBPh, 108 92 138 102 Pr,NBPh, 62 76 90 -14 Et,NBPh, -6 210 70 -228 Bu4NC10, 20 80 81 52 Pr4NC104 -25 -7 30 -30 Et,NCIO, -91 -80 12 -228 Bu,NI 12 112 130 -72 CuClO, * 4CH,CN -40 24 49 -8[Cu(DMPhen),]CIO, -150 230 90 -222 J. CHEM.SOC. FARADAY TRANS., 1994, VOL. 90 Table 2 Test of the additivity rule for limiting apparent molal com- negative or small K$* values for most of the ions in AN pressibilities in AN and BN at 298 K therefore indicate stronger solvation of ions in AN than in BN.K$'~O-~cm3 mol-' bar-' electrolyte pair AN BN The authors are grateful to the CSIR, New Delhi for a research grant under the research scheme 1( 122 1)/9 1 -EMR-I1Bu4NBPh4-Bu4NC104 88 57 and for the award of a senior research fellowship to J.S. and a Pr4NBPh4-Pr,NC10, 87 60 Et,NBPh,-Et,NCIO, 87 58 Research Associateship to V.A.Bu,NBPh,-Pr,NBPh, 46 48 Bu,NC10,-Pr,NC104 45 48 Bu,NBPh,-E t,N BPh4 114 68 BU~NCIO,-E~~NCIO~ 111 69 References Pr,NBPh,-Et,NBPh, 68 20 Pr,NCI0,-Et4NC104 66 18 1 I. Davidson, G. Perron and J. E. Desnoyers, Can. J. Chem., 1981, 59, 2212. 2 M. S. Bakshi, J. Singh, S. K. Bhullar, B. Kaur, S. C. Sharma and I. M. Joshi, Acoustica, 1992, 75, 292. Table 3 Limiting ionic apparent molal isentropic compressibilities 3 D. S. Gill, T. Kaur, H. Kaur, I. M. Joshi and J. Singh, J. Chem. of some ions in AN and BN at 298 K SOC.,Faraday Trans., 1993,89, 1737. 4 J. A. Riddick, W. B. Bunger and T. K. Sakano, Organic Solvents, cm3 mol-' bar-' Physical Properties and Methods of Purijication, Wiley Inter- ~g,/lO-~ science, New York, 4th edn., 1986. ion AN BN 5 D.S. Gill, K. S. Arora, J. Tewari and B. Singh, J. Chem. Soc., Faraday Trans. I, 1988,84,1729. Bu4N+ 48 62 6 D. S. Gill and B. Singh, J. Chem. Soc., Faraday Trans. I, 1988, Pr4N+ 3 11 84,4417. Et,N+ -63 -7 7 D. S. Gill, A. N. Sharma and H. Schneider, J. Chem. SOC., +cu -12 30 Faraday Trans. 1, 1982,78,465. Cu(DMPhen) -64 71 8 D. S. Gill, K. S. Arora, B. Singh, M. S. Bakshi and M. S. Ph,B -60 76 Chauhan, J. Chem. Soc., Faraday Trans., 1991,87, 1159. c10; -28 19 9 R. D. Lisi, S. Milioto and R. E. Verrall, J. Solution Chem., 1990, 1--36 68 19, 665. 10 F. Millero, J. Phys. Chem., 1971, 75, 280. 11 R. Zana, G. Perron and J. E. Desnoyers, J. Solution Chem., 1980, 9, 59. K$* values for the smaller Pr,N+ and Et,N+ ions are much 12 D. S. Gill and M. B. Sekhri, J. Chem. Soc., Faraday Trans. I,lower than that of Bu,N+ in AN and BN. For Ph,B-, the 1982, 78, 119. large Kgi value may also arise due to the change from a loose 13 D. S. Gill, J. Chem. SOC.,Faraday Trans. I, 1981,77, 751. structure into a more compact structure by the application of 14 D. S. Gill and A. N. Sharma, J. Chem. SOC., Faraday Trans. I, pressure from the sound wave. 1982, 78,475. Very small or negative tc$* values usually arise from enhanced structural effects due to the solvation of ions. The Paper 3/05018A; Received 18th August, 1993

ISSN:0956-5000    

DOI:10.1039/FT9949000579

出版商:RSC    

年代:1994

数据来源: RSC

摘要:

J. CHEM. SOC. FARADAY TRANS., 1994, 90(4),583-586 Transport and Compressibility Studies of Some Copper([) Perchlorates in Binary Mixtures of Benzonitrile and Acetonitrile Dip Singh Gill,* Rajinder Singh, Vazid AH, Jasbir Singh and Sharwan Kumar Rehani Department of Chemistry, Panjab University, Chandigarh-160014, India Transference numbers of CIO, and Cu+ ions in copper(i) perchlorate have been measured in the concentration range 0.024.22mol dm-3 in benzonitrile(BN)-acetonitrile(AN) mixtures at 298 K by the modified Hittorf method. The limiting transference numbers of Cu+ (t",) in various solvent systems were evaluated by the modified Longsworth method. Molar conductances, ultrasonic velocities and densities of CuCIO, * 4CH,CN and [Cu(DMPhen),]CIO, were measured at 298 K in BN-AN mixtures.The limiting molar conductances (Ao) of the electrolytes were obtained by analysis of the conductance data using the Shedlovsky equation. Limiting ionic conductances (2:) for Cu', [Cu(DMPhen),]+ and CIO, ions and hence their solvated radii (ri)were calculated. The variation in solvated radii of ions as a function of solvent composition showed no preferential solvation of any ion by AN or BN. From the ultrasonic velocity and density measurements, the isentropic compressibilities (K~) and the limiting apparent molal isentropic compressibilities (K:) for CuCIO, 4CH,CN and [Cu(DMPhen),JCIO, were evaluated. That K: values for the copper(i) salts became less negative or more posi- tive with increasing BN mole fraction indicating that both these copper(i) salts are structure makers in AN and that the structure-making tendency of the salts decreased with increasing BN mole fraction. Previously'v2 we reported transference number measurements of copper(1) perchlorate in AN and its binary mixtures.To the best of our knowledge, there have been no other trans- ference number measurements of copper(1) salts. Compress- ibility studies of copper(1) salts are also lacking. BN is another potential solvent which stabilises copper(1) salts. Copper@) perchlorate tetraacetonitrile (CuClO, 4CH3CN) is much more soluble in BN than in AN. In this paper we have measured the transference numbers, molar conductances and ultrasonic velocities of two copper(1) salts, CuC10, 4CH3CN and bis(2,9-dimethyl- 1,1O-phenanthroline)copper(r) perchlor-ate ([Cu(DMPhen),]ClO,) in BN and BN-AN mixtures.Experimental BN (Puriss, >99%GC, Fluka Chemica) and AN (99% from E.Merck, India) were further purified by the methods report- ed The density, viscosity and permittivity of the purified solvents were in good agreement with the literature values. CuClO, 4CH3CN and [Cu(DMPhen),]ClO, were pre-pared by the methods reported earlier.'*2*6 A modified Hittorf transference cell (75 cm3 capacity) with three compartments separated by well greased stopcocks was used in all transport-number measurements. The cathode and anode compartments of the cell were fitted with pure (99.9%) copper spirals which acted as electrodes. The electrodes were thoroughly cleaned with dilute nitric acid, then with distilled water and finally with dry acetone and were then dried.Weighed samples of CuClO, -4CH3CN were dissolved in the appropriate solvent or mixture to obtain solutions of the desired concentration. A direct current of 4 mA (measured accurately with a digital current meter as well as by a silver coulometer connected in series with the cell) was passed through the solution with the help of a current stabilizer (Hindustan Power Tronix Inc., New Delhi) for 4-6 h to bring about a measurable change in the Cu+ concentration. The solutions from the middle and the cathode compartments were analysed for Cu+ concentration by titration with aqueous cerium(rv) ammonium sulfate s~lution.~During analysis, the copper(1) solutions in pure AN or in the AN-rich region did not present any difficulty but in pure BN or the BN-rich region of the mixtures some difficulty arose as BN is not readily miscible with water.This difficulty was, however, overcome by adding an excess of AN to a solution of the copper(1) salts to be titrated and by shaking the solutions vigorously. The other details of measurements and the equa- tions used were as given earlier.8 The accuracy of the trans- ference number measurements was f1%. Conductances were measured at lo00 Hz with a calibrated digital conductivity meter. Details of the conductance mea- surements were reported previo~sly.~ The accuracy of con- ductance measurements was f0.2%.Ultrasonic velocities were measured at 2 MHz using an ultrasonic time intervalometer model UTI- 101 from Innova- tive Instruments, Hyderabad using a pulse-echo overlap technique. The absolute accuracy of the sound velocity mea- surements was 2 parts in lo4. Densities were measured using an Anton Paar Digital Densimeter (model 60) and a calibrated cell type 602 with a reproducibility of +O.OOOOl g cm-3. Results and Discussion Transference-number Measurements Transference numbers of ClO, and Cu' ions in copper(1) perchlorate were measured in the concentration range 0.02-0.22 mol dmP3 in BN-AN mixtures at 298 K by the modified Hittorf method. Limiting transference numbers of the Cu+ cation (t&+) were obtained by the modified" Longsworth method of extrapolation.The Longsworth function for Cu + (t;70,+)was calculated by the method reported by Kay and Dye" using the equations: t;ou+= tCu+ + (0.5 - tCu+)A,/AO (1) and A, = bC''2/(1 + ~d) (2) Where p = 82.487/q(~T)'/~,= 50.2916 C'/2/(~T)'/2K and 6 is the ion-size parameter, set equal to the Bjerrum critical dis- tance, 4 = e2/2&kT.The permittivity (E), viscosity (q) and density (p) for the BN-AN mixtures are reported in Table 1. A. values for CuClO, were obtained from conductance mea- surements. Plots of the Longsworth function (tpu+)us. con-centration (C) were linear (Fig. 1) and gave limiting Table 1 Permittivity, viscosity, density and ultrasonic velocity of BN-AN mixtures at 298 K BN(molYo) E VIP p/g c~r-~ u/m s-l 0 36.00 0.003 410 0.776 85 1280.8 11.4 33.96 0.004517 0.821 11 1300.8 25.5 32.40 0.005 388 0.865 41 1323.6 43.4 30.24 0.006 776 0.911 62 1352.0 67.2 27.80 0.008 63 1 0.954 43 1382.8 100 25.16 0.011 958 1.00034 1418.0 transference numbers of Cu+ cation (tg,+; Table 2) on extrapolation to infinite dilution by the least-squares method.The tg,+ values show a strong solvent composition depen- dence. They decrease significantly with increasing BN mole fraction in the mixture. Conductance Measurements Molar conductances of CuClO, -4CH3CN and [Cu(DMPhen),]ClO, were measured in the concentration range (1-60) x lop4 mol dm-3 at 298 K in BN-AN mix-tures. The limiting molar conductances (A,) and ion-association constants (K,) in all cases were determined by analysing the conductance data by the Shedlovsky equation' ' using a procedure reported previ~usly.'~.' The A, values for CuClO, and [Cu(DMPhen),]ClO, are report- ed in Table 2 along with some literature values for these salts.' 3-1 Using the tg,+ and A, values for CuClO, from Table 2, limiting ionic conductances (A:) for Cu' and ClO, ions have been calculated (Table 3) using the equations: Ao(C~C104)t&+ A&+ (3)= Ao(CuC104)= A:,+ + A&oz (4) Table 3 also lists the 1; values for [Cu(DMPhen),]+ in various BN-AN mixtures obtained by combining the A, values for [Cu(DMPhen),]ClO, from Table 2 with the corre- sponding A&oz values. Our A&+ and A&,, values (64.64 and 103.26 S cm2 mol-' in pure AN in Table 3) are in good agreement with the values (64.7 and 103.3 S cm2 mol-') reported by Yeager and Krat~chvil'~and Kay and co-workers.' g From the A: values of Table 3 the solvated radii (ri) for various ions have been calculated by using the eq~ation.'~ r.= -''' F2 + 0.0103~+ ry' 67rNqA9 J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 0.400 I I + 0;c 0.275 0.250 0 0.05 0.10 0.15 0.20 0.25 C/mol dm-3 Fig. 1 Longsworth function for Cu+ cation (t2u+)us. molar concen- tration of CuCIO, * 4CH3CN in BN-AN mixtures at 298 K. mol% BN: 0,O; 0, 100.11.4;A, 25.5; *, 43.4; x, 67.2; 0, Where F is the Faraday constant, N is Avogadro's number, q and E are the solvent viscosity and permittivity and ry is an adjustable parameter, taken as 0.085 nm for AN, BN and all BN-AN mixtures.The solvated radii for Cu+, [Cu(DMPhen),]+ and ClO, ions thus obtained are listed in Table 3. The ri values for Cu' and [Cu(DMPhen),] cations+ increase linearly with the BN mole fraction in the mixture. The ri value for ClO, does not show any variation with solvent composition. There is, therefore, no preferential solva- tion of Cu', [Cu(DMPhen),]+ or ClO, in AN, BN or AN-BN mixture^.^*'^^'^^^^ The increase in ri for Cu+ and [Cu(DMPhen),]' with BN mole fraction is simply due to the replacement of AN molecules by the more voluminous BN molecules in the solvation sphere of these ions with increas- ing mole fraction of BN.Ultrasonic Velocity Measurements Ultrasonic velocities (u) and densities (p) of binary mixtures of BN and AN as well as of CuC10,.4CH3CN and [Cu(DMPhen),]ClO, in these mixtures containing 0, 11.4, 25.5, 43.4, 67.2 and 100 mol% BN were measured in the con- centration range 0.01-0.15 mol kg- ' at 298 K. The ultrasonic velocities of the binary mixtures are reported in Table 1 while the ultrasonic velocities as a function of molality of the copper(r) salts are presented in Fig. 2. Fig. 2 shows that in both cases the ultrasonic velocities increase linearly with the molality of the electrolytes. Also the ultrasonic velocity Table 2 Limiting transference numbers of Cu+ cation in CuCIO, and A. and K, values for CuClO, and [Cu(DMPhen),]ClO, in BN-AN mixtures at 298 K CuClO, BN(mol%) Ao/S cm2 mol-'t:U f 0.385 167.9 (1 68.2)b (1 68.4)c (1 68.0)d (167.8)" 11.4 0.376 134.3 25.5 0.366 105.5 43.4 0.346 83.0 67.2 0.332 64.1 100 0.324 43.5 a Ref. 17; ref.13; ref. 14; ref. 15; " ref. 16. Cu(DMPhen),ClO, KJdm' mol-' Ao/S cm2 mol-' KJdm3 mol-' 17 155.6 (159.4)8 3 123.9 98.2 3 77.8 5 61.0 42.5 J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 Table 3 Limiting conductances and solvated radii (ri)in BN-AN mixtures at 298 K cu + [Cu( DMPhen),] c10;+ BN(mol%) Ip/S cm2 mol-' rJnm iLO/S cm2 mol -' rJnm Ip/S cm' mol- ' r,/nm 0 64.64 0.49 52.34 0.57 103.26 0.36 11.4 50.50 0.49 40.10 0.57 83.80 0.34 25.5 38.61 0.51 31.31 0.60 66.89 0.34 43.4 28.72 0.54 23.52 0.63 54.28 0.34 67.2 2 1.28 0.56 18.18 0.64 42.82 0.34 100 14.09 0.60 13.09 0.64 29.41 0.34 increases significantly with increasing BN mole fraction in the ent molal isentropic compressibilities, K# for both electrolytes mixture.were calculated using the equations given previously.21 The The isentropic compressibility (K~)of each electrolyte was limiting apparent molal isentropic compressibilities (K:) were calculated using the relation extrapolated from the linear plots of K# us. m1I2 by the least- squares method using the equation Ks = W2P) (6) K, = icg + A,m'I2 (7)The isentropic compressibilities for CuClO, .4CH,CN and [Cu(DMPhen),]ClO, decrease linearly with increasing Plots of K# us.m1'2are shown in Fig. 4. and the best extrapo- molality of the salts in BN-AN mixtures (Fig. 3). The appar- lated K: and A, values are reported in Table 4. There are no 90 1410 80 1360 70 601310 50 7 1260 I L E 0 0.05 0.10 0.15 0.20 0.25 2 40--. WI 90--. x1410 80 1360 70 60 1310 50 1260 0 0.02 0.04 0.06 0.08 0.10 40 0 0.02 0.04 0.06 0.08 0.10 m/mol kg-m/mol kg-' Fig. 2 Ultrasonic velocity us. molality for (a) CuCIO, .4CH3CN Fig. 3 Isentropic compressibility us. molality for (a) and (b) [Cu(DMPhen),]ClO, in BN-AN mixtures at 298 K. CuCIO, .4CH,CN and (b) [Cu(DMPhen), JCIO, in BN-AN mix-Symbols as in Fig. 1. tures at 298 K. Symbols as in Fig. 1. Table 4 Limiting apparent molal isentropic compressibilities and the slope of eqn.(7) for CuCIO, .4CH,CN and [Cu(DMPhen),]ClO, in BN-AN mixtures at 298 K CUCIO, * 4CH,CN [Cu(DMPhen),]CIO, BN(mol%) KP~O-~an3mol-' bar-' 104~~ K$'~O-~cm' mol-' bar-' 104A, 0 -40 24 -150 230 11.4 -8 -83 -56 -88 25.5 -1 -82 -29 -187 43.4 8 -39 -4 -160 67.2 10 -5 21 -221 100 49 -8 90 -222 0 -2 0 -40 -60 0.1 0.2 0.3 0.4 0.5 d1 o 50 F\ sz" 0 -50 -1 00 -1 50 L I I I I I 0.05 0.10 0.15 0.20 0.25 0.30 0.: m /2/mo11/2kg-1 /2 Fig. 4 Apparent molal isentropic compressibility us. square root of molality for (a) CuClO, * 4CH,CN and (b) [Cu(DMPhen),]CIO, in BN-AN mixtures at 298 K. Symbols as in Fig. 1. compressibility data on copper@ salts with which our K: values from Table 4 can be compared. The maximum uncer- tainty in K: values of Table 4, on the basis of our previous comparison of other salts, is estimated to be _+3 x lop4cm3 mol-' bar.-' The results indicate that both copper(1) salts are structure makers in AN and the structure-making effects decrease with increasing mole fraction of BN in BN-AN mixtures.The authors are grateful to the CSIR, New Delhi for a research grant under the research scheme 1( 1221)/91-EMR-II J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 and for the award of Senior Research Fellowship to J.S. and a Research Associateship to V.A. The authors thank Ranbir Singh for the computer analysis of the conductance data. References 1 D. S. Gill, K. S. Arora, J. Tewari and B.Singh, J. Chem. SOC., Faraday Trans., 1988,84,1729. 2 D. S. Gill, J. Tewari, G. Singh and M. S. Bakshi, J. Chem. SOC., Faraday Trans., 1991,87,1155. 3 D. S. Gill, J. Solution Chem., 1979,8, 691. 4 D. S. Gill, T. Kaur, H. Kaur, I. M. Joshi and J. Singh, J. Chem. SOC., Faraday Trans., 1993,89, 1737. 5 J. A. Riddick, W. B. Bunger and T. K. Sakano, Organic Solvents, Physical Properties and Methods of Purijication, Wiley Inter- science, New York, 4th edn., 1986. 6 B. J. Hathaway, D. G. Holah and J. D. Postlethwaite, J. Chem. SOC.,1961, 3215. 7 J. Bassett, R. C. Denny, G. H. Jeffery and J. Mendham, Vogel's Textbook of Quantitative Inorganic Analysis, Longman, London, 4th edn., 1978. 8 J. 0. Wear, C. V. McNully and E. S. Amis, J. Inorg. Nucl.Chem., 1961,18,48. 9 D. S. Gill, A. N. Sharma and H. Schneider, J. Chem. SOC., Faraday Trans. I, 1982,78,465. 10 R. L. Kay and J. L. Dye, Proc. Natl. Acad. Sci. USA, 1963,49, 5. 11 R. M. Fuoss and F. Accascina, Electrolytic Conductance, Inter-science, New York, 1959; R. M. Fuoss and T. Shedlovsky, J. Am. Chem. SOC., 1949,71,1496. 12 D. S. Gill, and M. B. Sekhri, J. Chem. SOC., Faraday Trans. 1, 1982,78, 119. 13 D. S. Gill and M. S. Chauhan, Z. Phys. Chem. NF, 1984, 140, 139. 14 H. L. Yeager and B. Kratochvil, J. Phys. Chem., 1969,73, 1963. 15 D. S. Gill and R. Nording, 2. Phys., Chem. NF, 1983,136, 117. 16 D. S. Gill, N. Kumari and M. S. Chauhan, J. Chem. SOC., Faraday Trans. I, 1985,81,687. 17 K. Miyoshi, J. Phys. Chem., 1972,76, 3029. 18 C. H. Springer, J. F. Coetzee and R. L. Kay, J. Phys. Chem., 1969, 73,471. 19 D. S. Gill, Electrochim. Acta, 1979,24,701; 1977,22,491. 20 D. S. Gill, S. Chauhan and M. S. Chauhan, Z. Phys. Chem. NF, 1986,150, 113. 21 J. Singh, T. Kaur, V. Ali and D. S. Gill, J. Chem. SOC., Faraday Trans., 1994,90,579. Paper 3/05019J; Received 18th September, 1993

ISSN:0956-5000    

DOI:10.1039/FT9949000583

出版商:RSC    

年代:1994

数据来源: RSC

摘要:

J. CHEM. SOC. FARADAY TRANS., 1994, 90(4), 587-590 Kinetic Model for Serum Albumin Adsorption :Experimental Verification Roger Kurrat, Jeremy J. Ramsdent" and Jiri E. Prenosil Department of Chemical Engineering, ETH-Zentrum, 8092Zurich, Switzerland When human serum albumin is adsorbed at a hydrated metal oxide surface, the amounts which can be desorbed by replacing the protein solution with pure solvent steadily decrease. A quantitative kinetic model involving two modes of adsorption, reversible and irreversible, is inferred, and tested by allowing the adsorption and desorption to take place on the surface of planar optical waveguides, and determining the surface coverage from the shift in the mode indices of guided waves due to the adsorption. The method allows the surface coverage to be determined to a precision of ca.1 ng cm-2 every 15-30 s. Excellent agreement with the model was obtained. Protein adsorption has captured the attention of researchers for many years, partly because of its essential technical and medical importance, and partly because of its many baming aspects.',2 There has been no shortage of plausible theoreti- cal models pr~posed,~,~ but many of them predict adsorption kinetics which differ slightly from one model to another, and to test them measurements of high precision are needed. Up until now, the experimental results have only been in qualit- ative agreement with model prediction^.^ A novel approach to this problem is to exploit the well known phase shift which occurs upon total internal reflexion at the boundary between two dielectrics of differing refractive indices (the same phenomenon which is the basis of the tech- niques of ellipsometry4 and reflectometry').In the case of a guided wave propagating in a thin optical waveguide, only discrete modes exist, whose eigenvalues N can be calculated from the opto-geometrical parameters of the For a four-layer waveguide, comprising support S, high refractive index film F, adsorbed adlayer A and cover medium C, the mode equation linking the refractive indices n,, n,, n, and n,, and thicknesses d, and d, -arctan[ (:)2pJ( -)I nf--NZ (z)2pJ(-arctan[ -)]nf--N2 ( 1) where p = 0 and 1 for the transverse electric (TE, containing the field components E,, H, and HJ) and transverse mag- netic (TM, containing the components H,, Ex and E,) modes, respectively, and m = 0, 1, ..., the mode number.The mea- surement of N for any one mode can be used to determine one unknown parameter. A layer of adsorbed protein mol- ecules is characterized by two parameters, its thickness d, and its refractive index n,; measurement of N for two modes allows both nA and d, to be determined by simultaneously solving eqn. (1). ~~ ~~ t Also at : Department of Biophysical Chemistry, Biozentrum, 4056 Basel, Switzerland. 1The z axis is perpendicular to the waveguide which lies in the x, y plane. The TE waves are so called because only the electric field has a transverse component. The adsorbed mass M per unit area of surface is, for a uniform adlayer of thickness d, and concentration within the layer C, 'O,' ' M = (CA -cb)d, (2) where cb is the bulk concentration, and is related to the frac- tional surface coverage, 8, by 8 = Ma/m (3) where a is the area occupied per molecule and m the mass of a single molecule. The refractive index of a protein solution depends linearly upon its concentration according to' nA = n, + cAdn/dc (4) where n, is the refractive index of the pure solvent, and the coeflicient dn/dc depends on the polarizability of the protein and has a quasi-universal value of 0.182 cm3 g-'.lo Combin-ing eqn. (2)-(4), and taking cA % cb ,yields 9= -kIdA m dn/dc (5) The mode indices N can be very conveniently determined if the waveguide incorporates a diffraction grating (period = A).Then, light from an external beam of wavelength I will be coupled into the waveguide provided its angle of incidence a onto the grating satisfies the conditiong*' N = nair sin a + lA/A (6) where nair is the refractive index of the external medium and 1 the diffraction order. Measurement of the angles at which the incoupled power is at a maximum (a series of discrete peaks are observed) allows N to be determined for all possible modes in the waveguide. Experimental Planar optical waveguides of TiSiO,, with a thickness of d, z 180 nm and refractive index nF z 1.8, with an incorpor- ated grating coupler (A = 416.15 nm) were obtained from AS1 AG, Zurich, Switzerland.These are monomode wave- guides which only allow the zeroth TE and TM modes to pro- pagate. They were pre-treated in hot concentrated Caro's acid for 30 min to remove any organic matter, rinsed exten- sively in doubly distilled water, and equilibrated overnight in buffer solution [lo mmol dm-3 "(2-hydroxyethyl) piperazine-N-ethanesulfonic acid-NaOH (HEPES) at pH 7.341. Human serum albumin (Calbiochem) was of high purity (>99%) and was used as received. A small silicone to voltmeter k-air bubbles \\measuring aperture Fig. 1 Measuring cuvette. The measuring aperture is the grating in the surface of the optical waveguide (‘chip’), onto which the external beam (He-Ne laser, A = 632.8 nm) is incident from below. The ther- mocouple for monitoring the temperature is located at the bottom of the tube for extracting the air bubbles, barely projecting into the cuvette proper.The cuvette has a semicircular cross section of 1.7 mm2, a radius R of 1 mm and is 8 mm long; the length x from the middle of the inlet tube to the middle of the measuring aperture (the grating coupler region) is 3.5 mm. rubber flow-through cuvette was sealed to the surface of the waveguide over the grating region such that the waveguide formed one wall of the cuvette (Fig. 1). Protein solutions of concentration cb were drawn through the cuvette using a peristaltic pump at a rate F of 0.56 mm3 s-‘. A second pump, operating at a quarter of the rate of the first, drew out material before the measuring area; the object of this was to remove any adventitious gas bubbles in the inlet stream which would disturb the guided modes.Temperature was monitored by a thermocouple inserted into this second outlet port. The measurements of a were carried out using an IOS-1 goniometer scanning device (AS1 AG, Zurich), which allows a to be determined with a resolution of 1.25 x rad, and hence N [eqn. (6)] to be determined to an accuracy of +2 x The light source was a linearly polarized He-Ne laser, I, = 632.8 nm, oriented such that the plane of polariza- tion made an angle of 45” with the plane of reflexion; hence the TE and TM modes were excited with equal intensities. The Model Under the experimental conditions used, the Reynolds number is ca.1. Flow is laminar, and since the diffusion coef- ficient of proteins is very small, the Peclet number is large, of the order 1O00, and hence the hydrodynamic boundary layer is about ten times thicker than the diffusion boundary Under these conditions the flux I to the surface is given by an expression of the form DcdS, where 6 is the diffu- sion boundary layer thickness. If every molecule arriving at the surface were immediately and irreversibly adsorbed, dM/dt would be equal to I. However, some of the adsorbed protein may be desorbed by flushing the cuvette with pure buffer solution, although the proportion desorbed tends to decrease as the surface fills up. Therefore, a reversible process must also be taken into account. Let us define rate constants k, and kd for adsorption and desorption, respectively, and in addition a rate constant ki for J.CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 irreversible adsorption. The adsorbed molecules are divided into two types, reversibly and irreversibly adsorbed, with coverages M, and Mi, respectively, such that M = M, + Mi (7) k, and ki may be a priori and, as will be seen, are a posteriori different from one another. Now we define a layer directly above the surface, in which the concentration of protein is c, (Fig. 2). At first glance, the total net rate of adsorption (i.e. the rate of change of the experimentally measured quantity M) is, according to the scheme of Fig. 2, dM/dt = (k, + ki)cS$ -kd M, (8) where $ denotes the fraction of the surface available for adsorption.Continuity in this layer, i.e. dc$dt = 0, means that I + k, M, = (k, + k&, $ (9) where I is now the net flux to the layer: I = (Cb-c,)D/6 (10) I may be eliminated between eqn. (9) and (10) to give an expression for c, The thickness of the diffusion boundary layer is given by” 6 = (3/2)(D~R/u)”~ (12) In this equation x and R depend on the dimensions of the flow cell and are defined in the legend to Fig. 1, and U is the maximum velocity of the fluid in the cell, given by U = 3F/(2A) (13) where A is the cross-sectional area of the cell. According to the Langmuir model, $ = 1 -8. Since pro- teins being adsorbed at random positions on the surface leave gaps between themselves too small to enable other proteins to be adsorbed (multilayer formation can be excluded because the adsorption always tends to a definite plateau), this expression does not accurately describe protein adsorp- ~~ layer directly above P DA the surface desorptionI,, ladsorption ?sorption~ ,surface I ka Cb -cs Mr kil kd Mi Fig.2 Reaction scheme for serum albumin adsorption. The distance of the layer directly above the surface from the surface is much smaller than 6. Reversibly adsorbed proteins are drawn resting on one of their apices, and irreversibly adsorbed molecules resting on a side. The lower part shows the rate constants. J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 0.24 I 1 NI 6 0,xs 0 0 10000 Fig.3 Typical experimental curves showing M us. time, t ( x). The solid line shows eqn. (6) fitted to the experimental data. The calcu- lated curves using the fitted parameters are: M,(t), (----) [from eqn. (16)]; Mi@),(....) [from eqn. (17)]. Conditions were c,, = 82.1 pg m-', F = 0.56 mm' s-', T = 25 f0.5"C. The arrows labelled a and b indicate, respectively, when protein flow began, and when it was replaced by pure buffer flow. tion, as has been verified by recent experiments.'3,'4 Instead, taking the excluded area properly into account,' 5,16 4 = 1 -4e + (qyn)e2 -[176/(3~2) -40/(.,/3)]e3 + o(e4) (14) where 0(04) are terms of order O4.t Results Fig. 3 shows a typical experimental curve (M us. t). After 5800 s adsorption, the protein solution was replaced with pure buffer.In order to fit the curves to eqn. (8), the following procedure was adopted :each parameter was optimized (using the least-squares criterion) in turn, holding the others con- stant. These optima defined a new search direction, in which the procedure was repeated, and so on until the desired M I b r Fig. 4 Sketch of the influence of the four fitting parameters on the shape of simulated M(t)curves Table I Fitted parameters collected from experiments, with condi- tions as in Fig. 1 k,/1OP6 cm s-l kds -I ki/10-6 cm s-l aa 3.5 -t 1.9 130 k60 1.9 f0.8 15+4 a is calculated from the fitting parameter m/a using a molecular weight of HSA = 65000 g mol-'. '* t Dickman et a1." have calculated the fourth-order term. degree of accuracy was attained.The adjustable parameters used in fitting were k,, ki, and the ratios k$ka and m/a. m is a constant and equal to 1.08 x g.18 Each of these parameters has a distinctly perceptible effect on the shape of a simulated kinetic curve (Fig. 4), and a rough optimum could be found by fitting by eye. Typically, each curve com- prised ca. 100 points; for the four adjustable parameters con- ditions for fitting were therefore very stringent. The parameters D and 6 were fixed: D was taken to be 3.261 x cm2 s-', measured for bovine serum albumin,' which is hydrodynamically almost identical to human serum albumin and 6 was calculated from eqn. (12) and (13) using the cell dimensions given in the legend to Fig.1. Several plausible variants on the model depicted in Fig. 2 were examined initially, but had to be rejected because it was imposible to fit them to the data. For example, addition of a parameter describing the conversion of reversibly adsorbed to irreversibly adsorbed protein was found to have a negligi-ble effect on the goodness of fit, and no good fit was possible if the distinct pathway of irreversible adsorption (parameter ki) was replaced by the conversion of M, to Mi on the surface. Furthermore, it was impossible to fit the simple expression for the rate of desorption, kdM, included in eqn. (8), to the data. In order to obtain satisfactory fits, an inverse square root dependence on the free area had to be included,t i.e.(dM/dt)dcsorption = -kd Mr/J4 (15) The complete model is therefore dMr/dt = ka c, 4 -kd Mr/J4 and dMJdt = kit,$ (17) the experimentally measured quantity being M=Mi+M,, with cb D/6 + kd Mr/J$c, = (ka + kiM + D/S The fitted parameters are collected in Table 1. Discussion From these results the following points could be deduced: (1) The significance of a: The hydrodynamically similar bovine serum albumin is a prolate ellipsoid with major and minor axes of 14 and 4 nm, respectively.20 The experimen- tally determined area a of 14 nm2 therefore implies a vertical orientation of human serum albumin at the solid/liquid inter- face. (2) There are separate pathways for reversible and irrevers- ible adsorption. Although the protein is quite symmetrical regarding its shape chemically it is not at all, and different orientations will bring different amino acids into contact with the surface.These will have different energy bar- riers for adsorption. Our choice of just two rate constants for adsorption, ka and ki, implies that the differences have been subsumed into two distinct groups, which also have different rate constants for desorption, kd and 0 respectively.$ (3) There is no interconversion of reversibly and irrevers- ibly adsorbed forms at the surface. When a particle is desorbed, it will initially attempt to be readsorbed in the t As well as -4, various other powers of 4 were tried but none were satisfactory. $ We intend to correlate these rate constants with the surface chemistry of the protein as soon as the crystallographic coordinates become available.590 0.0006 r I I tn \ NI \ \5 \2--. \ \ \313 \ '. 0 0 6000 Fig. 5 Rate of adsorption (---) and rate of desorption (-) us. time. Both approach zero asymptotically. vicinity of its former adsorption site.21*22 If during its excur- sion into the solution its orientation is randomized, such that irreversible adsorption is also possible, desorption + adsorption would provide a pathway equivalent to the con- version from reversibly to irreversibly adsorbed at the surface (by means of a conformational or orientational change of the adsorbed protein). Good fits could only be obtained by means of separate steps for irreversible and reversible adsorp- tion, implying that the correlation time for rotational motion of the particles is considerably longer than for lateral motion.(4) The rate of desorption is inversely proportional to the square root of the free area available for adsorption, an unex- pected finding. The root free area available for adsorption is proportional to a characteristic linear dimension of the gaps between adsorbed molecules. Hence the protein is desorbed more readily from a crowded interface; the presence of neigh- bours either impedes the formation of hydrogen bonds with the substrate, or else neighbours repel each other electrostati- cally (the isoelectric point of human serum albumin is at about pH 4.5,23and the experiments were carried out at pH 8).Fig.5, showing how the rates of adsorption and desorption vary with time, illustrates a key feature in protein adsorption. As the space available at the surface becomes filled up, the rate of adsorption drops dramatically. Initially, a high pro- portion of the molecules are reversibly adsorbed, but they are slowly replaced by their irreversible congeners, and the rate of desorption drops slowly but inexorably to zero. This is of great significance in establishing a criterion for when to begin regeneration of fouled ultrafiltration membranes and other surfaces contaminated by adsorbed proteins.24 Summary Using planar optical waveguides as substrates for protein adsorption, the opto-geometric parameters of the adsorbed J.CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 protein layer can be calculated from the mode indices, which can rapidly and accurately be determined from the incoup- ling angle of an external light beam directed onto a diffrac- tion grating incorporated into the planar waveguides. From these opto-geometric parameters, the mass of adsorbed protein per unit area can be calculated to an accuracy of ca. 1 ng cm-2. A complete (packed, ordered) monolayer of human serum albumin would give M = 0.77 pg cm-2. The present method is therefore accurate to about 0.1YOof a monolayer. A simple kinetic model was found to give excellent agree- ment with the measured data. Two states of the adsorbed protein were required, irreversibly and reversibly adsorbed.These states do not interconvert on the surface, but have separate adsorption and desorption pathways, which are believed to correspond to orientations of the molecule distin- guished by the particular amino acids in contact with the surface. The rate of desorption was proportional not only to the mass of reversibly adsorbed protein, but also inversely proportional to the square root of the free area, implying an interesting type of lateral interaction between adsorbed pro- teins, in which neighbours promote their desorption. References 1 F. MacRitchie, Ado. Protein Chem., 1978,32, 283. 2 J. D. Andrade and V. Hladky, Adv. Polym. Sci., 1986,29, 1. 3 I. Lundstrom and H. Elwing, J. Colloid lnterface Sci., 1990, 136, 68.4 R. M. A. Azzam and N. M. Bashara, Ellipsometry and Polarized Light, North-Holland, Amsterdam, 1977. 5 P. Schaaf, Ph. Dejardin and A. Schmitt, Rev. Phys. Appl., 1985, 20, 631. 6 P. K. Tien, Rev. Mud. Phys., 1977,49, 361. 7 A. Ghatak and K. Thyagarajan, Optical Electronics, Cambridge University Press, Cambridge, 1989. 8 J. J. Ramsden, J. Statist. Phys., 1993, 73, 853. 9 K. Tiefenthaler and W. Lukosz, J. Opt. Soc. Am. B, 1989,6,209. 10 J. A. de Feijter, J. Benjamins and F. A. Veer, Biopolymers, 1978, 17, 1759. 11 W. Lukosz and K. Tiefenthaler,Sens. Actuators, 1988, 15, 273. 12 V. I. Levich, Physicochemical Hydrodynamics, Prentice Hall, Englewood Cliffs, 1962. 13 J. J. Ramsden, Phys. Rev. Lett., 1993, 71, 295. 14 J. J. Ramsden and J. E. Prenosil, J. Phys. Chem., submitted. 15 P. Schaaf and J. Talbot, Phys. Rev. Lett., 1989,62, 175. 16 P. Schaaf and J. Talbot, J. Chem. Phys., 1989,91,4401. 17 R. Dickman, J-S. Wang and I. Jensen, J. Chem. Phys., 1991, 94, 8252. 18 X. M. He and D. Carter, Nature (London), 1992,358,209. 19 M. L. Wagner and H. A. Scheraga, J. Phys. Chem., 1956, 60, 1066. 20 A. K. Wright and M. R. Thompson, Biophys. J., 1975,15, 137. 21 F. Rabinowitch, Trans. Faraday SOC.,1937,33, 1225. 22 B. Senger, P. Schaaf, J. C. Voegel, A. Johner, A. Schmitt and J. Talbot, J. Chem. Phys., 1992,97, 3813. 23 P. G. Righetti, G. Tudor and K. Ek, J. Chromatugr., 1981, 220, 115. 24 T. Hediger, Dissertation No. 7933, ETH, Zurich, 1985. Paper 3/05484E;Received 13th September, 1993

ISSN:0956-5000    

DOI:10.1039/FT9949000587

出版商:RSC    

年代:1994

数据来源: RSC

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J. CHEM. SOC. FARADAY TRANS., 1994, 90(4), 591-595 591 Primary Yields of Water Radiolysis in Concentrated Nitric Acid Solutions Ryuji Nagaishi, Pei-Yun Jiang, Yosuke Katsumura" and Kenkichi lshigure Department of Quantum Engineering and Systems Science, Faculty of Engineering, University of Tokyo, 73-1 Hongo, Bunkyo-ku, Tokyo 113,Japan The primary yields of water radiolysis have been evaluated as a function of nitric acid or nitrate concentration in the 6oCo y-radiolysis of nitric acid and 0.4 rnol dm-3 sulfuric acid with nitric acid or sodium nitrate solutions containing Ce'" and Cell', in the presence and absence of TI+. The radiolytic decomposition of water is signifi- cantly enhanced in these systems, from g,( -H,O) = 0.60 pmol J-' for 0.4 rnol dm-3 sulfuric acid-sodium nitrate solutions tog,( -H,O) = 0.76 pmol J-' for nitric acid solutions at high solute concentrations.In order to understand the radiolysis of concentrated nitric acid solutions, we have studied the direct action of radiation of nitric acid, which has been shown to give rise to NO,+ e,& and 0 + NO,.'.2 Since nitrate anion is a strong scav- enger of pre-hydrated (e;) and hydrated (e,) the radiolysis of water is significantly influenced in nitric acid solutions, which is evident from the fact that the radiolytic reduction of the Ce" to Ce"' is remarkably enhanced in nitric acid and acidic nitrate sol~tions.~-'~ However, there seems to have been no estimation of the primary yields of water radiolysis in concentrated nitric acid and acidic nitrate solu- tions, which are indispensable for a fuller understanding of the radiolytic processes.The radiolytic reduction of Ce" to Ce"' in nitric acid, 0.4 mol dmP3 sulfuric acid-nitric acid and 0.4 mol dmP3 sulfuric acid-sodium nitrate solutions in the presence and absence of Tl+ has been studied in the present work with 6oCo y-radiation. The primary yields of water radiolysis have been evaluated as a function of nitric acid or nitrate concentration for these systems. Experimental Nitric acid solutions, sodium nitrate and other chemicals were of the highest available purity and were used as sup- plied. The water was purified by distillation followed by fil- tration through a millipore system. Aerated solutions were irradiated at room temperature with a 3 kCi 6oCo y-source, which provided dose rates in the range 0.02-0.14 Gy s-l as determined by the Fricke dosimeter with G(Fe3') = 1.62 pmol J-'.The concentrations of Ce" were measured spec- trophotometrically using molar absorption coefficients deter- mined in the present study. The energy deposition was assumed to be proportional to the electron density and was corrected accordingly. Results and Discussion Molar Absorption Coefficients (E) of Ce" Stock solutions of 4 mmol dm', Ce(S04), in 0.4 mol dm-, sulfuric acid were prepared and the concentration of Ce" as determined on the basis of &(Ce", 320 nm) = 561 m2 mol-' at 298 K in 0.4 mol dmP3 sulfuric acid solution^.'^ Solutions of 0.4 mol dm-3 sulfuric acid with nitric acid or sodium nitrate containing 0.2-0.4 mmol dm- Ce(SO,), were pre- pared by diluting the stock solutions and the concentrations of Ce" were calculated.The absorbances of the solutions were measured and the molar absorption coefficients of Ce" were evaluated at selected wavelengths in the range 350-420 nm. The decadic &(Ce", 370 nm) values obtained are shown in Table 1 for 0.4 mol dm-3 sulfuric acid-nitric acid solu- tions and 0.4 mol dm-3 sulfuric acid-sodium nitrate solu-tions. For nitric acid solutions, a weighed amount of anhydrous (NH,),Ce(NO,), was dissolved into the solutions and the concentrations of Ce" (0.2-0.6 mmol dm-3) were calculated. The decadic &(Ce", 370 nm) values obtained are also shown in Table 1.The determination of Ce" in irradi- ated solutions was carried out at several wavelengths in the range 350-390 nm where only Ce" absorbs. The measure- ment and the molar absorption coefficients of Ce" are not affected by the presence of Ce"' and Tl', at least in the wave- length range 350-390 nm. Yields of Radiolytic Reduction of Ce'" to Ce"' Aerated nitric acid and 0.4 mol dm-3 sulfuric acid with nitric acid or sodium nitrate solutions containing Ce'" and Ce"' with or without T1' were irradiated and analysed for CeIV. The concentrations of Ce" were 0.4-0.6 mmol dmP3 before irradiation and the ratios of [Ce"']/[Ce"'] were kept between Table 1 Molar absorption coefficients (370 nm) of Ce" (m2mol-') 0 253 0.25 250 0.25 258 0.14 40.5 0.50 249 0.50 262 0.60 63.5 0.75 250 0.75 266 0.8 1 75.3 1.o 253 1.o 269 1.1 89.0 2.0 266 2.0 278 2.2 135 3.0 283 3.0 283 3.1 178 4.0 304 4.0 286 4.2 219 5.0 329 5.0 288 5.1 303 6.0 357 6.0 289 6.1 390 7.0 387 - 7.1 445 8.0 417 8.1 470 2.0 I 1 I 1.5 c I 7-0 --.E, h 1.0 20" I Y Q 0.50 0.0 0.0 2.0 4.0 6.0 8.0 [HNO, or NaNO,]/mol dm-, Fig.1 Yields of radiolytic reduction of Ce'" to Ce"' in the presence and absence of T1+ in nitric acid and 0.4 rnol dm-3 sulfuric acid with nitric acid or sodium nitrate solutions. Nitric acid: (0)without Tl';(a)5.0 mmol dm -TI 0.4 mol dm- ' sulfuric acid-nitric acid : (A)+ .without T1+; (A)1.0 mmol dm-j TI+.0.4 mol dm-3 sulfuric acid- sodium nitrate: (0)without Tl+;(.)1.0 mmol dm-3 TI+. J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 was neglected in the present study because of the small elec- tron fraction for sulfuric acid. H,O- e,;, H, H, , OH, H,O,, H' (1) gw(-H,O) = gw(ea; + H) + 2gdH2) = g,(OH) + 2gw(H,O,) (1) NO;(HNO,)-.,NO, + eJe,; + H') (2) gSl(-nitric acid) = gs1(N03)= gSl(ea;) (11) NO;-.,O+NO, (3) gs,(-NOi) = Ssz(N02)= gs20 (111) HNO, -., 0 + HNO, (4) d2(-HN03) = d2(HN02) = !&2(O) (IV) 0 + NO, + 0, + NO, (5) In the solutions containing Ce" and Ce"' without Tl', Ce"' is oxidized by OH and NO, and Ce" reduced by HNO, and H,02 as well as by e;/e,; and H uia formation of HNO, by the following reactions [( 1 1)-( 15)I.l' 6-' Assuming that 9' H, and H,O, are only formed from the radiolysis of water, the yields of G(-Ce") in the absence of TI', G,(-CeIV), can be expressed by eqn.(V) and the yields of oxygen G(0,) by eqn. (VIa) and (VIb), wheref, andf, are the electron fractions of water and nitric acid or nitrate anion calculated from their total (inner-sphere and valence) electrons, respectively, with f, +fw = 1, and a is the degree of dissociation of nitric acid. 3 and 3 during irradiation, with Ce"' being added prior to irradiation or introduced by pre-irradiation. The initial con- centration of T1' was 1.0 mmol dm-, for 0.4 mol dm-, sul- furic acid with nitric acid or sodium nitrate solutions and 5.0 mmol dmP3 for nitric acid solutions.The dose rates were ca. 0.02-0.075 Gy s-l in the presence of T1' and 0.06-0.14 Gy s-' in its absence. The yields of the reduction of CeW to Ce,"' G( -Ce"), were calculated from the [Ce'v]-dose curves with good linearity in the dose range studied, i.e. less than 330 Gy. The results are shown in Fig. 1. The radiolytic reduction of CetV to Ce'" in 0.4 mol dm-, sulfuric acid has been studied extensively and established as a standard dosimeter. l4 In nitric acid and 0.4 mol dm-, sulfuric acid with nitric acid or sodium nitrate solutions, the yields, G( -Cew), are remark- ably enhanced. Addition of T1' nearly doubles G( -Ce"'). Our results in 0.4 mol dm-, sulfuric acid with nitric acid or sodium nitrate solutions are in good agreement with those reported previously7-' whereas the present results in nitric acid solutions are appreciably higher than those reported by Bugaenko and Roschektaer" and Vladimirova et ~1.'~As shown in Fig.1, the three systems are slightly different from each other. Reaction Mechanism of Radiolytic Reduction of Ce'" to Ce"' The radiolysis of water generates e;, H, H,, OH, H,O, and H' as primary products and the radiolysis of nitric acid gives rise to NO, + e,; and 0 + Material balances were assumed for these processes. The yield, g(e,;), designates the sum of pre-hydrated and hydrated electrons. The oxygen atom formed in reactions (3) and (4) has been shown to be probably in the triplet state, O(,P), which reacts with nitrate anion to yield 0, + NO, by reaction (5)., The added metal ions, Ce", Ce"' and Tl', cannot compete for the oxygen atom owing to their low concentrations.The direct action of radiation on sulfuric acid gives rise to SO, + e,; ,15 which Ce"' + OH + H' + CetV+ H,O (6) OH + HNO, + H,O + NO, (7) Ce"' + NO, + CetV+ NO; (8) 2CetV+ HNO, + H,O + 2Ce"' + NO, + 3H+ (9) 2Ce" + H,O, + 2Ce"' + 0, + 2H' (10) e,; + H++ H (1 1) e,g(ep) + NO, + H,O + NO, + 20H-(12) H + NO, +NO, + OH-(13) NO, + NO2 = N204 (14) N,04 + H20 + HNO, + H' + NO, (15) G,(-Ce'V) =fwCs,(e,, + H) -s,(OH) + 2gw(H20,)1 +f,[ssl(e,,) -gs,(NO,)l + 4f,Cag,,(-NO;) + (1 -a)g;2(-HNO,)I (V) G(02) =fwgw(H202) +f,C.s,,(-NO,) + (1 -4s:2(-HNO3)I (VI4 G(02) = CG,(-CetV) + 2fwgw(H2)1/4 (VIb) In the solutions containing Ce" and Ce"' with Tl', the following reactions (16)-(18) occur additionally.The yields of G(-CerV) in the presence of Tl', G,(-Ce"), can be expressed by eqn. (VII) and the yields of oxygen G(0,) still by eqn. (VIa). Thus eqn. (VII1)-(XIII) can be derived from eqn. (V) and (VII) and the material balance eqn. (I)-(IV) giving the primary yields of water radiolysis obtainable experimentally on the basis of G,(-Ce"), G2(-Ce"), G(H2),gsl,gs2,and g:, , where G(H,) is the observed yield of H, in the solutions. Primary Yields of Water Radiolysis The yields of G(H2) in nitric acid and neutral sodium nitrate solutions are identical up to at least 8 mol dm-3.9 For the three systems investigated in the present study, the yields of G(H2) were assumed to be the same as those in nitric acid and sodium nitrate solutions.The yields of gsl, gs2and g12 have been evaluated as 0.50,0.16 and 0.21 pmol J-' for nitric acid solutions1i2 and were assumed to be the same for the three systems of the present study. The concentrations of molecular HNO, were neglected for 0.4 mol dmP3 sulfuric 0.80 0.60 r I -7 f 0.40 -3 Q 0.60 0.40 0.20 t 0.0 0.0 2.0 4.0 6.0 8.0 [HNO,]/mol dm-, Fig. 3 Primary yields of water radiolysis in 0.4 mol dm-3 sulfuric acid-nitric acid solutions. gw(-H20) (O),gw(ea; + H) (A), g,(OH) (B)9gw(H,O,) (a)and gw(H2) (0). acid-sodium nitrate solutions and were assumed to be the same as those in nitric acid solutions2' for 0.4 mol dmP3 sulfuric acid-nitric acid solutions.In the calculation of the electron fractions, the contribution of sulfuric acid was neglected and for 0.4 mol dm-, sulfuric acid-sodium nitrate solutions, the electron fractions were modified as f, = f(NO,)/Cf(H,O) +f(NO,)I and f, = 1 -A 3 where f(H2O) and f(N0;) are the real electron fractions of water and nitrate anion, which means that the energy absorbed by Na+, by hypothesis, was shared by water and nitrate anion pro- portionally to their electron fractions. Thus we can calculate the primary yields of water radiolysis by eqn. (1X)-(XIII).The results are shown in Fig. 2-4. The decomposition of water Om80r---0.60 c I 7 0.40 f 0.20 0.0 0.0 2.0 4.0 6.0 8.0 [NaN03] /mol dm-Fig.4 Primary yields of water radiolysis in 0.4 mol dm-3 sulfuric acid-nitrate solutions. gw(-H,O) (O),gw(e, + H) (A), g,(OH) (B), gw(H20,)(a)and gw(H,) (0). 594 0.30 0.20 I 7-0 -..5. h N E! tl Q 0.10 0.0 0.0 2.0 4.0 6.0 8.0 [HNO, or NaNO,]/mol drr3 Fig. 5 The yields of (40,)in nitric acid and 0.4 mol dm-3 sulfuric acid-sodium nitrate solutions containing CeIV and Ce"'. Predicted G(0,) by eqn. (VIa) in nitric acid: (0)dose rate 0.02-0.14 Gy s-', this work and 0.4 mol dmP3 sulfuric acid-sodium nitrate (0):dose rate 0.02-0.14 Gy s-', this work. Calculated (40,)as [G,(-Ce'") + 2fwgw(H,)]/4in nitric acid: (A) dose rate 0.0016 Gy s-'; (0)dose rate 1.6 Gy s-', ref.13 and 0.4 rnol dm-3 sulfuric acid-sodium nitrate (A), ref. 10. Measured G(0,) in nitric acid: (B) dose rate 0.0016 Gy s-'; (+) dose rate 1.6 Gy s-', ref. 13 and 0.4 mol dm-3 sulfuric acid +sodium nitrate (O),ref. 10. increases rapidly with nitric acid or nitrate concentration at lower concentrations whereas it increases slowly or reaches a constant value at higher concentrations with gw(-H,O) values as high as 0.76 pmol J-' in higher than ca. 1 mol dm-, nitric acid solutions. Apparently the three systems differ from each other, presumably, owing to the influence of acidity on spur reactions, the nuances of which are rather difficult to understand quantitatively at present if possible. A rational description might be as follows: the scavenging of e;/ea; by NO; /HNO, leads to the increases in g,(ea; + H) and gw(-H,O) because of the lower reactivities of the NO, species formed and its precursors whereas the competition for ea; (but not e,) by H+ suppresses this effect due to the high reactivity of the H atom formed.The scavenging of OH by undissociated HNO, leads to increases in g,(OH) and g,( -H,O) owing to the lower reactivity of the NO, radical formed. Thus it is expected that the increases in g,(ea; + H), g,(OH) and g,( -H,O) should be greatest in pure nitric acid solutions and least in the 0.4 mol dm-3 sulfuric acid-sodium nitrate solutions, which is in general agreement with the experiment results. The formation of 0, has been investigated and the values of G(0,) reported for nitric acid and 0.4 mol dm-3 sulfuric acid-sodium nitrate solutions.'0~'3 The values of (30,)cal-culated from eqn.(VIa) on the basis of the primary yields evaluated in the present study are compared with the experi- mentally measured values of G(0,) as shown in Fig. 5, which are in good agreement for 0.4 mol dm-, sulfuric acid-sodium nitrate solutions" and for nitric acid solutions at lower dose rate under more or less similar experimental conditions to the present study.', The lower G(0,) values in nitric acid solu- tions at higher dose rate (10-100 times higher than those in the present study)', are presumably due to the reaction of H,O, with HNOz2' because of the higher dose rate and lower initial concentration of [Ce'"], = 0.1 mmol dm-3.The J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 values of G(0,) can also be obtained from eqn. (VIb) which is valid even when reaction (19) occurs. As shown also in Fig. 5, the values of G(0,) derived as [G,(-Ce") + 2f,gW(H,)]/4 are in good agreement with the experimental values of G(0,) under various conditions for nitric acid solutions' and 0.4 mol dmP3 sulfuric acid-sodium nitrate solutions" even when eqn. (VIa) is not applicable. These facts are in quantitive agreement with the present reaction mechanism and the derived primary yields. HNO, + H,O, + H,O + NO, + 3H+ (19) Attempts have been made to evaluate the yields of G(H,O,) in nitric acid solutions.".22 In 0.4 rnol dm-3 sul- furic acid-5.0 mol dm-3 sodium nitrate solutions, the value of G(H,O,) has been deduced as 0.12 pmol J-'," in reason- able agreement with the value of G(H,O,) =f,gw(H2O2) = 0.14 pmol J-' derived in the present study.In nitric acid solutions containing sulfanilamide as a HNO, scavenger, G(H,O,) has been obtained as 0.083 (0.05),0.084 (O.l), 0.11 (OS), and 0.114 (1.0 rnol dm-, nitric acid) pmol J-',,, increasing with nitric acid concentration, which is fairly con- sistent with the trend predicted in the present study. Since sulfanilamide is a scavenger of OH and NO,,' the system containing sulfanilaide is presumably different from that of the present study, which may give different G(H,O,) values. The mechanism of water radiolysis has been established and the primary yields are well documented. In 0.4 mol dm-, sulfuric acid solutions, g,(ea; + H) = 0.383, gw(H2)= 0.041, g,(OH) = 0.30, gw(H,O,) = 0.083 and g,( -H,O) = 0.466 pmol J- ' as measured by the Fricke dosimeter.In the pulse radiolysis studies, the initial value of g,(OH) was found to be 0.611 pmol J-' at 200 PS,,~ and that of g,(ea;) to be 0.50 pmol J-' at 30 ps24 and 0.48 pmol J-' at 100 PS.,~ These initial yields have been generally accepted in computer simu- lation of spur reactions.26 The yield of water decomposition in nitric acid solutions is as high as 0.76 pmol J-', which may give considerably higher initial yields for OH and ea;, presumably, corresponding to earlier times. The yield of water decomposition (initial or at 10-2-10-' s) has been obtained as 0.90,,' 0.71,,* ca.0.8,29and 0.70,' pmol J- ' by therotical calculation, which is supported by the present results. This work was supported in part by a Grant-in-Aid for Scien- tific Research (No. B-04453163),from the Ministry of Educa- tion, Science and Culture of the Japanese Government. References 1 Y. Katsumura, P. Y. Jiang, R. Nagaishi, T. Oishi, K. Ishigure and Y. Yoshida, J. Phys. Chem., 1991,95,4435. 2 P. Y. Jiang, R. Nagaishi, T. Yotsuyanagi, Y. Katsumura and K. Ishigure, J. Chem. SOC.,Faraday Trans., 1994,90,93. 3 M. J. Bronskill, R. K. Wolff and J. W. Hunt, J. Chem. Phys., 1970,53,4201. 4 R. K. Wolff, M. J. Bronskill and J. W. Hunt, J. Chem. Phys., 1970,53,421 1.5 C. D. Jonah, J. R. Miller and M. S. Matheson, J. Phys. Chem., 1977,81,1618. 6 G. E. Challenger and B. J. Masters, J. Am. Chem. SOC.,1955, 77, 1063. 7 T. J. Sworski, J. Am. Chem. SOC.,1955,77,4689. 8 H. A. Mahlman, J. Phys. Chem., 1960,64, 1598. 9 H. A. Mahlman, J. Chem. Phys., 1961,35936. 10 H. A. Mahlman, J. Phys. Chem., 1963,67, 1466. 11 T. J. Sworski, R. W. Mathews and H. A. Mahlman, Ado. Chem. Ser., 1968,82, 164. 12 L. T. Bugaenko and B. M. Roshchektaev, High Energy Chem., 1971,5,424. 13 M. V. Vladimirova, A. A. Ryabova, I. A. Kulikov and A. S. Mil-ovanova, High Energy Chem., 1977,II, 130. J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 595 14 15 R. W. Matthews, Znt. J. Radiat. Zsot., 1982,33, 1159. P. Y. Jiang, Y. Katsumura, R.Nagaishi, M. Domae, K. Ishi-kawa, K. Ishigure and Y. Yoshida, J. Chem. SOC., Faraday Trans., 1992,88, 1653. 23 24 25 C. D. Jonah and J. R. Miller, J. Phys. Chem., 1977,81, 1974. T. Sumiyoshi, K. Tsugaru, T. Yamada and M. Katayama, Bull. Chem. SOC.Jpn., 1985,58,3073. C. D. Jonah, M. S. Matheson, J. R. Miller and E. J. Hart, J. 16 17 M. Gratzel, A. Henglein and S. Taniguchi, Ber. Bunsenges. Phys. Chem., 1969,73,646. G. V. Buxton, C. L. Greenstock, W. P. Helman and A. B. Ross, 26 Phys. Chem., 1976,80, 1267. J. A. La Verne and M. Pimblott, J. Phys. Chem., 1991, 95,3196, and references therein. 18 19 20 21 J. Phys. Chem. Ref. Data, 1988, 17, 513. P. Neta, R. E. Huie and A. B. Ross, J. 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ISSN:0956-5000    

DOI:10.1039/FT9949000591

出版商:RSC    

年代:1994

数据来源: RSC

摘要:

J. CHEM. SOC. FARADAY TRANS., 1994, 90(4), 597-604 Pulse Radiolysis Study of the Reactions of SO;-with some Substituted Benzenes in Aqueous Solution Getahun Merga, C. T. Aravindakumar and B. S. M. Rao* Department of Chemistry, University of Poona , Pune-4 1I 007, India H. Mohan and J. P. Mittal" Chemistry Division, Bhabha Atomic Research Centre, Bombay400 085,India The reactions of SO:-with several substituted benzenes having the general formula, C,H,-,X,Y (where X = F, CI or Br and Y = CH,, CH,CI, CHCI,, CF, or OCH,), have been investigated in aqueous solution by pulse radiolysis. The transient absorption spectra exhibit maxima at 315-330 nm and additional peaks at 270-290 nm with chlorotoluenes and weak peaks around 400 nm with chlorobenzene and 3-chlorofluorobenzene. Only in the case of 3-chloroanisole is the observed spectrum different, exhibiting two distinct peaks at 290 and 475 nm.The second-order rate constants for the reaction of SO:-range from about 10' for 2-chlorobenzotrifluoride to 10" dm3 mol-' s-' for 3-chloroanisole. It is concluded from the Hammett treatment (p' = -1.6) that the reaction mechanism involves both direct electron transfer and addition-elimination reactions. The intermediate radical cation is hydrolysed to give the corresponding 'OH adduct absorbing at 315-330 nm except in the case of 3-chloroanisole where it is stabilized. The formation of a benzyl-type radical by direct H abstraction by SO:-from the CH, group and/or deprotonation of the radical cation is an additional process whose extent is deter-mined by the relative position of the CH, group, the order being para > ortho x meta with monochlorotoluenes.The transient species absorbing around 400 nm is assigned to the phenoxyl-type radical. The differences in reaction mechanism between SO:-and 'OH attack are discussed. Radiation chemical studies on the reactions of primary radicals '-'of water and secondary radicals,'-14 derived from them, with benzene and its substituted derivatives in aqueous solution have contributed to an understanding of structure- reactivity relationships. In the radiolysis of aqueous solu- tions, the reactive species formed are solvated electrons (eJ, 'OH radicals and H atoms : ionizing H20 -eJ2.7), 'OH(2.8), 'H(0.6), H+(2.7), radiation H202(0.6), Hz(0.43) (1) The numbers in the parentheses represent the G values of the species, i.e.the number of molecules per 100 eV (one molecule per 100 eV = 1.036 x lo-' mol J-') of absorbed energy. Because of their high yields, the reactions of both 'OH and ea; with arenes have been widely studied (see ref. 15 and 16 for recent reviews). It is known that the reactivity of e,; with arenes is strongly influenced by the nature of the substituent as was seen" from the increasing rate of its reaction with high-electron-affinity substituents. In the case of halogenated aromatic compounds, the reaction may involve the n orbitals of the aromatic ring or of the substituent halogen The former, a pre- dominant process with fluorocompounds, results in the for- mation of the molecular anion, whereas other aromatic halogenocompounds give an aryl radical and a halide ion.The hydroxyl radical has been shown to react with substi- tuted benzenes first by addition to the benzene ring forming a 7t complex which immediately rearranges to a r~ complex. The different isomeric 'OH adducts',* formed from this rearrangement have absorption maxima in the range 310-340 nm. In our earlier study' on the reactions of the OH radical with substituted benzenes, the reaction rates for the forma- tion of *OH adducts were found to depend on the nature of the substituent. The oxidation of these adducts by K,Fe(CN), was reported6*' to depend on the nature of the substituent, and in the absence of an oxidant, they decay by second-order kinetics.The production of radical cations in aqueous solution by radiation chemical methodsg,' 1*1 relies on the use of oxidi-zing species such as SO;-, Clip, Bri-, TIZ+.Among these, the SO;-radical anion (Eo = 2.5 -3.1 V us. NHE' '*19) is com-e monly used and can be produced from the reactions of ea& and H [reaction (2)] with K2S20, or from 'OH with H2S04.20921 S,O;-+ eJH) -+SO;-+ SOt-(HSO,) (2) Laser flash photolysisZ2 of S,O;-[either by 193 or 248 nm (reaction 3)] has also been shown to be a clean source of so;-: hvs,o;--2so;-(3) ChemicalZ3 and thermal24 methods to produce SO;-were used to elucidate the reaction mechanism by product analysis. The SO:-radical anion reacts with several aromatic com- pounds to give hydroxycyclohexadienyl radicals whose for- mation was explained" either by the addition of SO;-to the benzene ring followed by hydrolysis or by a direct electron transfer from the ring to SO:-followed by hydration of the resulting radical cation.Substituted benzenes of the type C,H, -,X,Y (where X = F, C1 or Br, and Y = CH,, CH,CI, CHCI,, CF, or OCH,) contain different electron-donating and withdrawing groups and it is, therefore, of interest to investigate the reac- tions of SO;-with these systems as they form an ideal class of compounds for structure-reactivity studies. Furthermore, such an investigation would reveal the differences in reaction mechanism between 'OH and SO;-attack.Experimental Preparation of Solutions The substituted benzenes obtained from Fluka were of high purity (>98%) and were used without any further purifi- cation. Other chemicals used were of analytical grade. The solutions were prepared in water purified by the Millipore Milli-Q system and the solutions containing S,Oi-and tert- butyl alcohol were saturated with N, prior to the dissolution of the solute to avoid its volatilization during degassing. Irradiation The reaction of SO:-was studied in N,-saturated aqueous solutions containing K2S208(1.5 x rnol drnp3), solutes (10-4-10-3 mol dm-3) and 0.2 mol dm-3 tert-butyl alcohol. Since the rate constants for the reaction of 'OH radical with the solutes, as determined by us earlier,' are in the range (2- 9) x lo9 dm3 mol-' s-', the reactivity (k[S] = lo8 s-') of OH radical with tert-butyl alcohol (koH+Bu,OH= 5 x lo8 dm3 mol-' s-') is higher at least by an order of magnitude even in the case where the rate constant for the reaction of the OH radical with the solute is the highest.Furthermore, our recent work3 show that the reactivity of el with the compounds used in this study (lo6 s-I), under our experimental condi- tions, is lower by two orders of magnitude than its reactivity with S,Oi-(10' s-'). Therefore, more than 90% of the 'OH radicals will be scavenged by tert-butyl alcohol while e,; reacts quantitatively with S,Oi -to produce the SO;-radical ion [reaction (2)]. The rate of reaction of an H atom with SO:-(k = 2.5 x lo7 dm3 mol-' s-') is lower than that of its reac- tion with benzeneI7 and benzene derivatives (k z 9 x 108 dm3 mol-' s-').Under the conditions employed in this study, only a fraction of H (ca. 30%) is expected to yield SO:-via reaction (2). Thus, G(SO:-) was estimated to be 3.3 molecules per 100 eV [=G(e,i) at high S20i-c~ncentration,~+0.3G(H)]. The contribution of the H atom reaction with the solute, being negligible, is not taken into account. Pulse radiolysis experiments were carried out using high- energy 7 MeV electron pulses (pulse width 50 ns) from a linear accelerator at BARC, Bombay and the details of the set-up are publishedz6 elsewhere. A KSCN dosimeter was used in the optical pulse radiolysis using G~500= 21500 dm3 mol-' cm-' per 100 eV for the transientz7 (SCN);-. The dose per pulse, depending upon the pulse width, was in the range 10-20 Gy.-"I J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 The decay of SO:-in the absence and in the presence of various solutes was monitored at 460 nm. The transient absorption as a function of time was recorded on a storage oscilloscope interfaced to a computer for kinetic analysis.". Results and Discussion Absorption Spectra of the Transients The transient absorption spectra were measured from pulse radiolysis of the solutions containing the solute (1O-j mol dm-3), K,S,08 (1.5 x lo-' mol dm-3) and tert-butyl alcohol (0.2 mol dmP3) by monitoring the absorbance in the range 250-570 nm. The maximum absorbance of the tran- sients was reached within about 2 ps for systems whose rate constants are 2 lo9 dm3 mol-' s-'.Since the SO:-radical anion reacts with tert-butyl alcohol with a half-life of about 5 ps (vide infra) at the concentration employed in our experi- ments, appropriate corrections were applied to the transient absorption spectra with systems having k d 9 x 10' dm3 mol-I s-'. The values of the absorption maxima of the tran- sient species formed are compiled in Table 1. Toluene and Chlorobenzene The transient absorption spectra with the two mono-substituted compounds, toluene and chlorobenzene, are more or less similar with I,,, centred at 315 and 325 nm, respec- tively (Fig. 1). The spectrum in the case of chlorobenzene, in addition, exhibits a weak peak around 400 nm with a shoul- der around 270 nm. As the parent compound absorbs below 280 nm, the shoulder observed may not be characteristic of any intermediate product.The absorbance at 325 nm decays by about 20% over a period of 18 ps after the pulse but the absorbance at 400 nm still persists. Transient absorption spectra were also obtained for the reaction of 'OH with toluene and chlorobenzene by pulsing N,O-saturated solutions of these compounds. These spectra were more or less similar (A,,, x 325 nm) to those obtained in the reaction of SO;-with both systems, though the inten- sities are different in the case of chlorobenzene. Fig. l(a) shows both the spectra obtained in the reaction of 'OH and SO:-with chlorobenzene. Since G(SO:-) = 3.3, the yield of 30 20 10 C r L250 350 450 E 3 0 300 350 400 450 500 550 t 0 Ilnm l/n rn Fig.I Time-resolved transient absorption spectra obtained from the reaction of SO:-with (a)chlorobenzene and (b) toluene: (0)imme-diately, (x) 2 p,(a)10 ps after the pulse. Transient absorption spectrum measured at 2ps after the pulse from the reaction of 'OH with chlorobenzene (A),normalized to G('0H) = G(SOi-) = 3.3 for direct comparison. Dose per pulse =20 Gy, pH 5.5. J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 'OH is normalized to 3.3 [from G('0H) = 5.61 for direct comparison with the spectrum observed in the reaction of SO;-. In the reaction of 'OH radical, the absorbance observed at 325 nm is higher with no absorption at 400 nm in contrast to the spectrum measured in the case of SO;-attack.The molar absorptivity of the 'OH adducts of chloro- benzene at 325 nm is determined to be 4600 dm3 mol-' cm-'. The transient species formed at 315 and 325 nm in the reaction of SO;-with toluene and chlorobenzene, respec- tively, may be attributed to their 'OH adducts. This assign- ment is in accord with the results obtained from our present study as well as with those previously rep~rted.~ The formation of 'OH adducts in the reaction of SO;-with benzene derivatives was explained' by both heterolysis [reaction (7), Scheme 11 of the initially formed SO;-adducts [reaction (5)] or via the hydration [reaction(lO)] of the ini- tially formed radical cation [reaction (4)].However, a clear- cut distinction between the two types of mechanisms cannot be made as the life-time of the SO;-adduct to benzene was reported" to be <lo0 ns; even in the case of benzonitrile, a compound with a strongly electron-withdrawing -CN group, the observed rate constant l5 for hydrolysis is >5 x lo6 s-'. The heterolysis rate was found" to increase with the introduction of a methyl substituent in the case of cyclohexene and ally1 alcohol which indicates an S,1 type hydrolysis of the SO;-adduct. Based on a pf value of -2.4 found with some substituted benzenes and benzoates, Neta et al. suggested" that the reac- tion mechanism proceeds by electron transfer from the benzene ring to SO;-. Our p+ value of -1.6 (vide infra) seems to suggest that the reaction proceeds both by direct electron transfer and addition-elimination processes Y Scheme 1 Generalized mechanism for the reaction of SO:-with substituted benzenes of the type C,H,-,X,Y (where X = F, C1 or Br and Y = CH,, CH,CI, CHCl,, CF, or OCH,).The intermediate radical cation is formed by both direct electron transfer [reaction (4)] and S,1 hydrolysis [reaction (9)] of the SO;-adduct formed via reaction (5). The radical cation is hydrolysed to give the correspond- ing 'OH adduct [reaction (lo)] except in the case of 3-chloroanisole. The formation of 'OH adduct radical via reaction (7) may not be likely under the experimental conditions employed. The phenoxyl radical is formed by reaction (8) only to a minor extent in the case of chlorobenzene and 3-chlorofluorobenzene.Reactions (6)and (12) are additional channels for mono- and di-chlorotoluenes. [reactions (4) and (5)]. Since reaction (7) is expected to be base-catalysed, it may not be predominant under our experi- mental conditions (pH 5.5). It can be concluded, therefore, that the formation of 'OH adducts occurs via the hydration of a radical cation formed directly [reaction (4)] or from the SN1 hydroysis [reaction (9)] of the SO:-adduct. The intermediate with a weak peak around 400 nm in the case of chlorobenzene is, possibly, due to the formation of a phenoxyl radical by the elimination of HSO, from the radical adduct [reaction (@I.In another recent study2' involving the reaction of *OH with 3-chloroanisole under acidic conditions, we have observed absorption maxima at 315 and 400 nm which are attributed to the corresponding 'OH adduct and phenoxyl radicals, respectively.It is also reported that phenoxy13' and trichlor~phenoxyl~radicals absorb around 400-430 nm. The alternative possibility of the formation of the radical cation [reaction (9)] of chloroben- zene may be unlikely as its khydtation of benzene. ' should be as high as that Dihalogenobenzenes and 3-Chloroanisole The transient absorption spectrum with 3-chloro-fluorobenzene exhibits two well defined peaks at 315 and 400 nm (Fig. 2). As compared with the spectrum with chloroben- zene, there was an increase in absorbance at 400 nm while the intensity at 3 15 nm decreased, indicating that the elimination of HSO, is more predominant.However, only a band with A,,, at 330, with no additional peak at 400 nm, was noticed in the case of 1,3-dibromobenzene suggesting that SO:-reacts quantitatively leading to the formation of the *OH adduct. The corrected transient absorption spectrum is shown in Fig. 2. The spectrum with 3-chloroanisole, having a strong electron-donating -OCH, group, shows two intense peaks at 290 and 475 nm and a shoulder around 325 nm within 2 ps after the pulse (Fig. 2). The nature of the spectrum clearly demonstrates the formation of the radical cation as observed ~earlier' [E nm ~= 7240~ and E~~~ nm = 3800 dm3 mol-cm -'3 with methoxy- and dimethoxy-benzenes.The radical cation in this case is stabilized owing to the presence of the electron-donating -OCH, group and its hydration [reaction (lo)] is, therefore, unlikely. The calculated molar absorpti- vities are 5600 and 3000 dm3 mol-' cm-' at 290 and 475 nm, respectively. The decay of the absorbance at these wave- lengths was considerable as can be seen from the time- resolved spectra measured at 2 and 16 ps after the pulse. In order to obtain information on the nature of the decay, absorption traces were taken at 290 nm on a longer timescale (100 and 500 ps) both at low (<OSkrad per pulse) and high (1.5 krad per pulse) doses. The fitting of the data in the former case, where the bimolecular decay is negligible, was not satisfactory owing to the poor signal-to-noise radio.The decay of the radical cation at high dose was monitored for three different concentrations (0.05,0.1 and 0.2 mol dm-3) of tert-butyl alcohol. The t1,2 obtained from absorption traces (inset of Fig. 2) is about 13 ps and such a decay indicates a first-order process. We cannot, however, conclude from our data whether the decay of the radical cation is due to its reaction with tert-butyl alcohol or S20i-, or to its rearoma- tization to form the phenyl-type radical [reaction (1 l)]. Mono- and Di-halogenotoluenes The spectral intensities of compounds having both -CH, and -C1 substituents (monochlorotoluenes) seem to depend on their relative positions. The spectra with 2-and 3-chlorotoluenes are more or less similar to that observed in the case of toluene with prominent peaks at 320 and 325 nm, J.CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 40 30 30 s 20 9)0 C C e-2j20 nE m X X m02 10 10 0 I I I I 0 2 3 300 350 400 450 500 550 2 I 300 350 400 450 500 550 Ilnm I/nm Fig. 2 Transient absorption spectra obtained from the reaction of SO;-with (a) 3-chloroanisole, (x) 2 ps and (A) 16 ps after the pulse; (6) 2-chloro-6-fluorotoluene (@), 3-chlorofluorobenzene (m), at 2 ps and 1,3-dibromobenzene (V)at 4 ps after the pulse, corrected for the com- petition reaction of SO>-with tert-butyl alcohol (see text). Dose per pulse =20 Gy, pH 5.5. The inset shows the trace for the decay at 290 nm in the case of 3-chloroanisole at a dose of 15 Gy per pulse and 0.2 mol dm-3 tert-butyl alcohol.respectively. In addition, a weak peak around 260-270 nm is observed with these systems which is considerably enhanced in the case of 4-chlorotoluene. This peak is assigned to the benzyl-type radical which is confirmed from the spectrum obtained by pulse radiolysis of N,O-saturated solutions of 4-chlorotoluene (lo-, mol dm-3) at pH 13, where the 'OH radical exists as 0'-(OHeH' + 0.-,pK = 11.9).,' It is known4 that 0'-reacts preferentially with the methyl group by abstraction of an H atom. The resulting spectrum from the reaction of 0'-with 4-chlorotoluene with absorption at 270 and 315 nm is nearly identical with that reported' in the case of 2-chlorotoluene. Both these spectra have similar fea- tures to those obtained from the reaction of SO;-with 4- chlorotoluene (Fig.3). The formation of benzyl-type radicals in high yield in the case of 4-chlorotoluene is due to H abstraction from the 30 A/nm CH, group [reaction (6)]. It is difficult to explain the ineffi- cient abstraction from 2-and 3-chlorotoluenes, though the steric effect may hinder the abstraction process in the case of 2-chlorotoluene. The formation of a benzyl-type radical, partly by deprotonation of the radical cation [reaction (12)] cannot, however, be ruled out. Pulse conductivity data with variation in pH are needed to obtain more quantitative infor- mation regarding the contribution of each of the two pro- cesses, as our results with absorption spectroscopy at pH 3, 5.5 and 9.6 are not conclusive.The contribution of SO:-leading to the formation of the benzyl-type radical is estimated to be about 10 and 40% for 3-and 4-chlorotoluenes, respectively, based on the reported4 E~~~~~ value of 14000 dm3 mol-' cm-' for the benzyl radical. Our spectral results therefore suggest that the order of the yields of 'OH adducts with monochlorotoluenes is 7c 60 50 9)uC 40 0 9 x 30 m2 20 10 0 250 300 350 400 450 500 550 Alnm Fig. 3 Transient absorption spectra obtained from the reaction of SO;-with (a) 2-chlorotoluene (FJ) and 3-chlorotoluene ( x) (corrected for the competition reaction of SO:-with tert-butyl alcohol). (6) Transient absorption spectra obtained from the reaction of SO;-(FJ) and 0'-( x ) with Cchlorotoluene.All spectra are recorded at 2 ps after the pulse. Dose per pulse =20 Gy. J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 30 20 Cm e s X n 0-1c C 1 *I 250 300 350 400 450 500 550 0 300 350 400 450 500 550 l/nm Ifnm Fig. 4 Transient absorption spectra obtained from the reaction of SO;-with (a) 3,4dichlorotoluene (@) and 2,6-dichlorotoluene (m)both 10 ps after the pulse. Both the spectra are corrected for the competition reaction of SO:-with tert-butyl alcohol. (b) 3-Bromotoluene ( x ) 2 ps after the pulse. Dose per pulse =20 Gy, pH 5.5. meta > ortho > para. Our preliminary data32 on product analysis by the HPLC technique from the reaction of SO;-with monochlorotoluenes in the presence of Fe(CN)i -under steady-state conditions have shown that the yields of the pro- ducts formed from benzyl-type intermediates are very small in the case of 2- and 3-chlorotoluenes as compared with 4- chlorotoluene.The extent of the formation of the benzyl-type radical seems to depend on the position of -C1 on the benzene ring. This is evident from the differences in the observed intensities at 270 nm in the spectra with 3,4- and 2,6-dichlorotoluenes. The spectrum with 3,4-dichlorotoluene is more or less similar to that observed with 4-chlorotoluene, whereas the spectrum with 2,6-dichlorotoluene resembles those of 2-and 3-chlorotoluene (Fig. 4). One interesting observation of the absorption spectra from mono- and di-chlorotoluenes is that, when the C1 atom is at position 4 relative to the CH, group (e.g.4-chlorotoluene, 3,4-dichlorotoluene), the formation of the benzyl-type inter- mediate by abstraction of an H atom from the -CH, group is more favoured. The effect of the nature of the substituent on the transient spectra can be seen from a comparison of the spectra of 2- chloro-6-fluorotoluene and 2,6-dichlorotoluene (A,,, = 270, 330, 410 nm) with those of 1,3-dibromobenzene, 3-bromo- toluene and 3-chlorofluorobenzene. The spectrum of 2-chloro-6-fluorotoluene has two overlapping peaks at 300 and 330 nm and two additional small peaks at 390 and 490 nm (Fig. 2). Only peaks (A,,, = 320-330 nm) corresponding to their 'OH adducts [reaction (lo)) are formed with 1,3-dibro- mobenzene and 3-bromotoluene (Fig.4), whereas an addi- tional peak at 410 nm is noticed, as mentioned earlier, in the case of 3-chlorofluorobenzene. a-Fluoro-and a-Chloro-toluenes The spectra with compounds (3,4,a-trichlorotoluene,2-chlorobenzylchloride, 2,6,a,a-tetrachlorotoluene and 2-chlorobenzotrifluoride) having halogen substituents both on the benzene ring and on the side group, have similar features with peaks corresponding to 'OH adducts and weak maxima above 400 nm. Kinetics for the Reactions of SO;-Radicals Decay of SO;-Rates for the reaction of SO;-with substituted benzenes may be determined from the decay of its absorption at 460 nm. N,-saturated S20i-solutions (1.5 x mol dm-j) con- taining 0.2 mol dm-, tert-butyl alcohol were pulse radiolysed (10-20 Gy per pulse) to monitor the decay of SO;-in the absence of solute.A first-order decay of SO;-with tIl2 = 5 ps was observed [(b)of inset of Fig. 5). The possible reactions for the decay of SO;-in the absence of any solute are: (i) its self-reaction, so;-+ so;-+ s,o;-(13) k = 7.6 x lo8 dm3 mol-' s-' ,21 1.6 1.2 time 1 ps -4 c +-....IIu) 02 0.8 -2 -rc 0.4 0 0.2 0.4 0.6 0.8 1.0 [4-chlorotoluene]/l 0-3 mol dw3 Fig. 5 k,, as a function of [4-chlorotoluene]. Inset shows the decay of SO:-at 460 nm, (a) with and (b) without bchlorotoluene mol dm-'). Dose per pulse =20 Gy, pH 5.5. (ii) its reaction with S,Oi-.or any impurity, X-(e.g.C1-) in the solution, SO;-+ S,O;-(X-) --+ SO:-+ S,Og-o(') (14) k(S0;-+ S,Oi-) = 6.6 x lo5 dm3 mol-' s-',,' (iii)and its reaction with the 'OH radical scavenger, tert-butyl alcohol, SO4-+ (CH,),COH -,HSO, + 'CH,(CH,),COH (15) k = 9 x lo5 dm3 mol-' s-'.~~ The tl,, for the decay of SO;-due to reaction (14) under our experimental conditions ([S,Oi-] = 1.5 x lo-, mol dm-3) should be at least 70 ps. The first half-life for the second- order decay of SO;-by reaction (13) is calculated to be about 200 ps considering G(SO;-) = 3.3 and the dose per pulse of 20 Gy. The observed decay of SO;-cannot be ascribed to its reaction with the impurity, Cl-, which is present only to an extent of 0.005% in the K2S208 used by us. Thus, reaction (15) seems to be the major pathway for the observed decay of SO;-in the absence of solute.A value of 1200 dm3 mol-' cm-' was obtained for the molar absorptivity of SO;-at 460 nm measured relative to an assumed G~500nm = 21 500 dm3 mol-' cm-' for (SCN);- and G(SO;-) values of 1100-1600 dm3 mol- ' cm-' for SO;-measured at 440-450 using either pulse radiolysis or flash photoysis tech- niques. Evaluation of Second-order Rate Constants The second-order rate constants for the reaction of SO;-with the substituted benzenes chosen in this study are deter- mined from the least-squares fit of the plots of kobs us. solute concentration. The pseudo-first-order rate constant of SO;-with tert-butyl alcohol (kobs = 1.7 x lo5 s-') is taken as the Y intercept and a typical plot is shown in Fig.5 in the case of Cchlorotoluene. The values of the rate constants obtained with different solutes are summarized in Table 1. The highest second-order rate constant, among the investi- gated compounds, is observed with 3-chloroanisole (k = 1.2 x 10" dm3 mol-' s-'). The rate constant in this = 3.3. Earlier ~~rker~~~*~~-~~ have obtained E s-'3 owing to the deactivation of the benzene ring for SO;-attack by electron-withdrawing groups such as -CH,Cl, -CHCl, and -CF3 . The rate constants for the addition of the 'OH radical to the different solutes used in this study were reported8 by us to be in the range (2-9) x lo9 dm3 mol-' s-', whereas the values for SO;-attack are in the range (108-10'o) dm3 mol-' s-' indicating that the rate of the latter reaction is more influenced by the nature of the substituent.Such an observation was also made by Neta et al." who report rate constant values varying from 5 x lo9 for anisole to <lo6 dm3 mol-' s-l for nitrobenzene. Eflect of Structure on Reactivity by the Hammett Treatment It is known that the SO;-radical anion behaves as an elec- trophile and any substituent effect can be quantitatively cor- J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 case is determined from the growth of the absorbance at 290 nm, instead of that at 460 nm, to avoid interference from the intermediate transient product which also absorbs at 460 nm. The high reactivity of SO;-with 3-chloroanisole, as in the case of dimethoxybenzene~,~ can be attributed to direct elec- tron transfer [reaction (411 rather than to formation of an intermediate adduct to the benzene ring [reaction (5)].The rate constants for the reaction of SO;-with chloro- benzene, toluene, 2-, 3-, and 4-chlorotoluenes, 3-bromo- toluene, 3,4-dichlorotoluene, 3-chlorofluorobenzene and 2- chloro-6-fluorotoluene are 2lo9 dm3 mol-s-'. The second-order rate constants for these systems are determined from at least five different concentrations of the solute C(0.2- 1) x mol dm-3] and their accuracy is within +lo%. A comparison of the rate constants for the reaction of SO;-with chlorobenzene (k = 1.5 x lo9 dm3 mol-' s-') and benzene" (k = 3 x lo9 dm3 mol-' s-') shows that the reac- tion rate is affected by the -C1 group.The rate constants for 2-~hlorobenzotrifluoride,3,4,a-tri-chlorotoluene, 2,6,a,a-tetrachlorotolueneand 2-chloro-ben-zylchloride (whose rates are lower) were measured with relatively high concentrations C(O.6-1) x mol dm-3] ; however, the concentration range is restricted to < mol dm-3 of the solute owing to the solubility limit. These solutes have apparently low reactivity [k = (1-7) x 10' dm3 mol-' Table 1 Bimolecular rate constants (10' dm3 rno1-l s-') obtained in the reaction of OH and SO;-with some substituted benzenes bimolecular rate constant probablesite of substrate OH" so;- absorption maxima/nm UCJll attack chlorobenzene -1.5 270, 325, - 400 toluene -3.1 315,400 -0.31 3-chlorofluorobenzene 4.8 1.o 315, 410 - 3-chloroanisole 9.3 12.0 290,475 -0.58 2-chlorotoluene 6.5 1.7 260, 320 -0.17 3-chlorot oluene 3.5 0.9 260, 325 -0.11 Cchlorotoluene 5.5 1.1 270, 315 -0.31 3,4-dichlorotoluene 2,6-dichlorotoluene 1.7 -0.9 0.4 270, 315 270, 330 -0.11 +0.23 2-chloro-6-fluorotoluene 4.2 1.2 410 300, 330 +0.06 3,4,a-trichlorotoluene 2-chlorobenzyl chloride 2,6,a,a-tetrachlorotoluene 2.5 4.1 4.9 0.2 0.7 0.2 390,490 290, 330 325,460 290, 320 +0.34 +0.10 - 410 2-chlorobenzotrifluoride -0.1 315, 390 - 3-bromotoluene 4.9 1.7 320 -0.10 1,3-dibromobenzene -0.5 330 - " Taken from ref.8. J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 related by the Hammett equation. However, each of the substituted benzenes used in our study contains a halogen atom is one of the substituents and an electron-donating (-OCH,, CH,) or an electron-withdrawing (-CF,, -CHCl, ,-CHCl,) group at different positions.A similar procedure, as in the case of 'OH attack,8 involv- ing the algebraic sum (acal)of the Hammett constant^^^-^' (a; or a:) for para and meta substituents and the Taft constant42 (a*)for the ortho substituent was used to deter- mine the more probable site for SO;-attack where Q,,, is expected to be minimum. The minimum aCa1values for differ- ent compounds are shown in Table 1 along with the positions for the more probable site for SO:-attack. The rate constants for diffusion-controlled reactions are not normally expected to obey the Hammett relationship as was seen6 from the levelling in log k us.acalplots for values of log k z9.5 in the oxidation of hydroxycyclohexadienyl rad- icals by IrC1;- and Fe(CN);-. Since the log k values of the compounds chosen in this study are <9.5 except for 3- chloroanisole, the plateau is not apparent. In the case of 3- chloroanisole (k = 1.2 x 10" dm3 mol-' s-'), the selection of the site of attack must occur within the encounter complex as was suggested43 for the reaction of 'OH with biphenyl (k = 1.04 x 10" dm3 mol-' s-') and phenol (10" dm3 mol-' s-I). The Hammett-type plot shown in Fig. 6 with a p+ values of -1.6 has a good linear correlation indicating that a similar reaction mechanism is operative in all cases.Our p+ value of -1.6 is in between the values reported" for 'OH (p' = -0.5) and SO;-(p+ = -2.4) with substituted ben- zenes. Therefore, we suggest that the reaction proceeds by both direct electron transfer as well as addition4imination channels. Though H abstraction is noticed with 4-chlorotoluene and 3,4-dichlorotoluene, they are included in the Hammett plot as the abstraction process is not exclusive. It has been rep~rted~,~~,~~ that the 'OH radical behaves as a weak electrophile and reacts by addition to the aromatic ring with low selectivity (p' = -0.5). The Hammett treat- ment of our data shows that a better correlation is obtained in the case of SO:-reaction than for 'OH attack. This implies that the positions of the probable site of attack (Table 1) may have greater influence on the reaction rate in the former case.However, only from determination of the product distribution of the oxidation of different isomeric 'OH adducts under steady-state conditions can one obtain direct evidence regarding the probable sites of attack predict- lo.Or' -9.5 rn -2 9.0 -8.5 8.01 I I I -0.6 -0.4 -0.2 0 0.2 0.4 OG3l Fig. 6 Hammett plot for the reaction of SO:-with substituted benzenes. 1, 3-Chloroanisole; 2, toluene; 3, Cchlorotoluene; 4, 2-chlorotoluene; 5, 3-chlorotoluene and 3,4-dichlorotoluene; 6, 3-bromotoluene; 7, 2-chloro-6-fluorotoluene; 8, 2-chlorobenzyl-chloride ;9,2,6-dichlorotoluene; 10, 3,4,a-trichlorotoluene. ed by the gCalvalues. Such a product analysis with isomers of monochlorotoluenes is in progress.Conclusion The Hammett analysis and spectral features of the transients formed in the reaction of SO;-with substituted benzenes reveal that the reaction mechanism involves both direct elec- tron transfer and addition4imination processes, eventually leading to the formation of the corresponding hydroxy- cyclohexadienyl radicals. Only with 3-chloroanisole is the formation of a radical cation observed. H abstraction from the CH, group is an additional reaction channel, especially in the case of 4-chlorotoluene and 3,4-chlorotoluene. The extent of each of these processes is dependent on the nature and the relative position of the substituents. The reaction rates [k = (108-10'o) dm3 mol-' s-'1 in the case of SO:-attack are more influenced by the substituent than those of 'OH attack.This is also reflected in the magni- tude of the Hammett constants obtained with SO;-(p' = -1.6) and 'OH (p' = -0.5) and in the differences in the transient absorption spectra obtained with these two reacting species. This suggests that SO;-is a stronger electro- phile and is more selective in its reaction. This study demon- strates that substituted benzenes of the type C,H, -,X,Y (where X = F, Cl or Br and Y = CH, , CH2Cl, CHCl,, CF, or OCH,) are an interesting class of compounds for investi- gation of structure-reactivity relationships by radiation chemical techniques. The authors thank Dr. R.M. Iyer of the BARC and Prof. M.S.Wadia of the University of Poona for their interest in this work. They also thank the referees for their suggestions, particularly regarding the interpretation of the phenoxyl radical absorption spectra, and the DAE for supporting this work. C.T.A. is thankful to the CSIR for the award of a senior research fellowship. References 1 J. Lichtsheidl and N. Getoff, Znt. J. Radiat. Phys. Chem., 1976, 8, 661. 2 S. Gordon, K. H. Schmidt and E. J. Hart, J. Phys. Chem., 1977, 81, 104. 3 H. Mohan, M. Mudaliar, B. S. M. Rao and J. P. Mittal, In?. J. Radiat. Phys. Chem., 1992,40, 513. 4 K. Sehested, H. Corfitzen, H. C. Christensen and E. J. Hart, J. Phys. Chem., 1975,79,310. 5 S. Steenken and N. V. Raghavan, J. Phys. Chem., 1979,83,3191. 6 G.V. Buxton, J. R. Langan and J. R. L. Smith, J. Phys. Chem., 1986,90,6309. 7 P. B. Draper, M. A. Fox, E. Pelizetti and N. Serpone, J. Phys. Chem., 1989,93, 1938. 8 H. Mohan, M. Mudaliar, C. T. Aravindakumar, B. S. M. Rao and J. P. Mittal, J. Chem. SOC., Perkin Trans. 2, 1991, 1387. 9 P. O'Neill, S. Steenken and D. Schulte-Frohlinde, J. Phys. Chem., 1975,79,2773. 10 P. Neta, V. Madhavan, H. Zemel and R. W. Fessenden, J. Am. Chem. SOC.,1977,99,163. 11 S. Steenken, P. O'Neill and D. Schulte-Frohlinde, J. Phys. Chem., 1977,81,26. 12 V. Jagannadham and S. Steenken, J. Am. Chem. SOC., 1988, 110, 2188. 13 Q. Sun, G. N. R. Tripathi and R. H. Schuler, J. Phys. Chem., 1990,94,6273. 14 C. von Sonntag, The Chemical Basis of Radiation Biology, Taylor and Francis, London, 1987, p.40. 15 (a)S. Steenken, J. Chem. SOC., Faraday Trans. I, 1987, 87, 113; (b) S. Steenken, in Free Radicals in Synthesis and Biology, ed. E. Minisci, Nato AS1 Series C-260, Kluwer Academic, Dordrecht, 1989, p. 213. 604 J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 16 P. Neta and A. Harriman, Photoinduced Electron Transfer, ed. 30 N. V. Raghavan and S. Steenken, J. Am. Chem. SOC., 1980, 102, 17 M. A. Fox and M. Chanon, Elsevier, Amsterdam, 1988, p. 110. G. V. Buxton, C. L. Greenstock, W. P. Helman and A. B. Ross, 31 3495. J. W. T. Spinks and R. J. Woods, Introduction to Radiation 18 19 J. Phys. Chem. Re5 Data, 1988, 17, 513. C. T. S. Lian and J. P. Mittal, Radiat. Phys. Chem., 1984,24,209. L. Eberson, Adu.Phys. Org. Chem., 1982,18,79. 32 33 Chemistry, Wiley, New York, 1976, p. 259. G. Merga and B. S. M.Rao, 1993, unpublished. E. Hayon, A. Treinin and J. Wilf, J. Am. Chem. SOC., 1972,94,47. 20 B. Lesigene, C. Ferradini and J. Pucheault, J. Phys. Chem., 1972, 34 W. Roebke, M. Renz and A. Henglein, lnt. J. Radiat. Phys. 76, 3676. Chem., 1969, 1, 39. 21 P. Y. Jiang, Y. Katsumura, R. Nagaishi, M. Domae, K. Ishikawa and K. Ishigure, J. Chem. Soc., Faraday Trans., 1992,88,1653. 35 36 E. Hayon and J. J. McGarvey, J. Phys. Chem., 1976,71,1472. L. Dogliotti and E. Hayon, J. Phys. Chem., 1967,71,2511. 22 W. J. McElroy and S. J. Waygood, J. Chem. SOC., Faraday 37 V. E. Heckel, A. Henglein and G. Beck, Ber. Bunsenges. Phys. Trans., 1990,86, 2557. Chem., 1966,70,149. 23 F. H. Quina and D. G. Whitten, J. Am. Chem. SOC., 1975, 97, 38 0. P. Chawla and R. W. Fessenden, J. Phys. Chem., 1975, 79, 1603. 2693. 24 M. E. Snook and G. A. Hamilton, J. Am. Chem. SOC., 1974, W, 39 P. Neta and L. M. Dorfman, Adv. Chem. Ser., 1968,81,222. 860. 40 M. K. Eberhardt, J. Phys. Chem., 1975,79, 1930. 25 T. I. Balkas, J. H. Fendler and R. H. Schuler, J. Phys. Chem., 41 J. March, Advanced Organic Chemistry Reactions, Mechanism 26 1970, 74, 4497. S. N. Guha, P. N. Moorthy, K. Kishore, D. B. Naik and K. N. and Structure, McGraw Hill, Kogakusha Ltd., New York, 1968, p. 241. Rao, Proc. Indian Acad. Sci., 1987,99, 261. 42 J. Hine, Physical Organic Chemistry, McGraw Hill, New York, 27 E. M. Fielden, The Study of Fast Processes and Transient 1962, p. 97. Species by Electron Pulse Radiolysis, ed. J. H. Baxendale and F. 43 X. Chen and R. H. Schyler, J. Phys. Chem., 1993,97,421. Busi, Reidel, Boston, 1984, p. 59. 44 P. Neta and R. H. Schuler, J. Am. Chem. SOC.,1972,94, 1056. 28 M. S. Panajkar, P. N. Moorty and N. D. Shirke, BARC Rep., 1988, 1410. 29 M. Mohan, Ph. D. Thesis, University of Poona, 1993. Paper 3/05999E; Received 7th October, 1993

ISSN:0956-5000    

DOI:10.1039/FT9949000597

出版商:RSC    

年代:1994

数据来源: RSC

摘要:

J. CHEM. SOC. FARADAY TRANS., 1994, 90(4), 605-608 Dual-cylinder Microelectrodes Part 2.t-Steady-state Generator and Collector Electrode Currents B. J. Seddon,*$ Chang Fa Wang, Wenfeng Peng§ and Xueji Zhang Department of Chemistry, Wuhan University, Wuhan 430072,People’s Republic of China The feedback and collection diffusion current properties of platinum dual-cylinder microelectrodes have been investigated using bipotentiostatic chronoamperometry. The study focuses on the influence of the interelectrode spacing on the steady-state response at independently controlled generator and collector microcylinders. A series of such devices were fabricated onto polyethylene terephthalate (PET) strips using a microscope manipu- lation method. Micro-manipulation techniques allowed dual-cylinder devices to be constructed using 25 pm platinum wire with gap spacings from 3.5 to 188 pm and cylinder lengths 250-500 pm. Recycling efficiencies, expressed as an iJiQ ratio, observed with these feedback dual-microelectrodes could be as high as 83% for closely spaced microcylinders. Electrochemical measurements further indicate that a convergence of the gener- ator and collector currents would follow as the interelectrode distance diminishes below the pm scale where the strength of the recycling flux field at the device is high enough to minimise bulk solution exchange.The response of the new amperometric device is discussed with reference to the relevant theoretical stationary flux model proposed for dual-cylinder microelectrodes. Amperometric dual microelectrodes function according to interelectrode diffusion processes which bring about a rapid and amplified steady-state current response at two closely positioned microelectrodes.Microelectrode devices such as these can be operated in both feedback and collection modes by independent control and measurement of individ- ual electrode potentials and currents. In collection experi- ments, where a collector is potentiostatted at a specified constant potential, the electro-generated species diffuse out from the generator electrode, traverse a microscopic gap and then undergo further redox reactions at a neighbouring col- lector electrode. The measured collection currents may then be utilised to bring improvements to analytical selectivity by monitoring the generated form of a redox analyte, hence overcoming background effects suffered by the generator elec- trode and indeed complicating homogeneous reactions.As a feedback device, however, it is necessary to sustain a high cyclic flux of some electroactive redox couple in the electro- lyte region between the microscopic anode and cathode. In both cases, feedback and collection micro-devices provide a constant current output in techniques such as pulse ampero- metry. Amperometric signals are measured on the nA scale and possess exceptionally low collector background currents which allows sensitive measurements on sub-micromolar analyte concentrations. Further, the onset of a steady-state current, i.e.response time, for these devices is of the order of a few seconds or less, with amperometric amplification over that of single microelectrode devices. Diffusion effects at dual microelectrodes provide us with new possibilities for the development of electroanalytical and sensor techniques over the conventional amperometric microelectrodes whose response is governed primarily by bulk solution mass trans- Dual microelectrodes like the interdigitated arrays (IDA), parallel dual cylinders (PDC) and feedback disc have been cited in a number of analytical problems over recent years, from end-point detectors in amperometric titrimetry,’ t Part 1: Ref. 14. Present address: Laboratoire d’Electrochimie, Institut de Chirnie Physique 3, Ecole Polytechnique Federale de Lausanne, CH-1015 Lausanne, Switzerland.0 Present address: Laboratory of Electroanalytical Chemistry, Changchun Institute of Applied Chemistry, Chinese Academy of Sci-ences, People’s Republic of China. to homogeneous immunoassay,6 collection sensing of cate- cholamines and stripping ~oltammetry.~ The experimental variables of importance in feedback and collection electrode and sensor studies are associated with the absolute and relative values of steady-state generator and col- lector currents. Collection current and efficiency terms are most frequently applied. Consequently, much of the charac- terisation work on multiple-band microelectrodes has relied on terminology from earlier diffusion-onvection experi-ment~.”~’~In this case the collection efficiency is the ratio of collector to generator currents. This term suitably marks the ability of a dual-electrode device to detect electro-generated species.The quantity recovered at the collector over trans- port losses is dependent on chemical and hydrodynamic con- ditions and electrode dimensions. The definition is quite satisfactory for rotating ringdisc experiments and dual-electrode flow cells since the generator electrode process remains independent of the collector electrode reaction owing to the macroscopic size of the electrodes and gap, as well as the directionality of solution flow. The description is also suitable for microelectrode systems where no feedback effects are concurrent.Alternatively, in instances where inter-electrode feedback diffusion is present, the generator current becomes modulated by a reverse transport process, i.e. the feedback flux. The observed collector current is then not in response to an independent generator process and hence under such conditions it is inappropriate to use the term col- lection efficiency in the former sense. For microelectrode feed- back devices generally, the ratio of collector current to generator current, as given in eqn. (I), is a measure of the diffusion recycling flux between generator and collector elec- trodes over any reaction or bulk exchange fluxes. Bearing this in mind, we have assigned the collector current: generator current ratio in these measurements at dual-cylinder micro- electrodes as a recycling efficiency which is sensitive to the physico-chemical conditions concerning the device and elec- trolyte solution.N = iJig Dual-cylinder microelectrodes were recently developed as an alternative geometry to laminate dual-band microelec- trodes,’0-’2 and as a more accessible, inexpensive and robust electrode system relative to the microlithographic IDAs.’ A dual-cylinder device is composed of a parallel arrangement of two cylindrical electrodes with microscopic dimensions in the cylinder diameter and the electrode spacing (for analytical applications <25 pm). By contrast, the electrode length is an order of magnitude or more larger; generally >250 pm but typically < 1 mm.The microelectrode devices examined in the present study are described as geometrically symmetrical in terms of cylinder diameter and spatial arrangement as shown in Fig. 1.14 Dual-cylinder microelectrodes possess certain features important to the design of feedback and col- lection experiments. The microscope fabrication scheme offers an extremely versatile method for the incorporation of multiple cylinders with a wide range of electrode and fibre materials, such as carbon fibre, metal micro-wires, chemically modified wires and coated polymer fibres.I3 Moreover, the well defined microcylinder geometry allows for a detailed characterisation of electrochemical systems by semi-analytical and numerical approaches to mass transport and coupled reaction flux problems.._ L d Fe(CN):-Fe(CN),& Fig. 1 (a) Illustration of the dual-cylinder microelectrode showing the device dimensions: cylinder length L, radius T and the gap spacing w, where w is the shortest separation distance between a pair of parallel cylindrical electrodes and is given by 2(d -r). (b) Electro-chemical recycling scheme for the hexacyanoferrate(lII/II) couple at cylindrical generator and collector microelectrodes. J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 A theoretical description of the diffusion current for a reversible redox couple at a dual-cylinder microelectrode was recently put forward in Part 1, where the effects of solution conductivity, convection and diffusion into bulk solution were neg1e~ted.l~ In this treatment it was assumed that a steady-state flux field is established between the cylindrical electrodes which culminated in a constant current of equal magnitude at the generator and collector, that is, no redox analyte losses into the bulk electrolyte.This current, which is dependent on the physical properties of the analyte system and the device dimensions, is re-stated in eqn. (2) below : i = nnFDCbL/cosh-'[(w/2r) + 13 (2) where i is the steady-state current delivered at the elec- trodes, with Cband D,the bulk solution concentration and diffusion coefficient of the electroactive species. The number of electrons involved in the electrode reaction is n, with L,r and w being the device dimensions as described in Fig.1. In this work an experimental study of the steady-state currents at the generator and collector of dual-cylinder devices is pre- sented for the case of Fe(CN);-/Fe(CN):-recycling in aqueous KCl electrolyte. The observations are discussed in the context of the theory for diffusion currents at these novel feed back microelectrodes. Experimental Chemicals Chemical reagents used were of analytical standard. Pot- assium hexacyanoferrate(rr1) and potassium chloride were purchased from Shengyang and Chantou Chemical Factories, China. All solutions were prepared with doubly distilled deionised water. The analytical solutions were purged with dry nitrogen for 10 min before electrochemical measurements were made. Electrode fabrication materials included, 'Melinex' polyethylene terephthalate (PET) kindly donated by ICI Films Ltd, platinum wire purchased from Goodfel- lows, and epoxy-resin from RS Components.Dual-cylinder Microelectrodes Dual-cylinder microelectrodes were fabricated by a novel microscope manipulation technique involving a slit-mounting and epoxy-sealant procedure. The dual electrodes are sup- ported on a PET strip for convenience of handling and oper- ation. Details concerning the construction and electrochemical characterisation of dual-cylinder devices have been given elsewhere. 59 Device dimensions were determined by scanning electron microscopy or with an optical micro- scope fitted with micrometer verniers. It was found that the length of the micro-wire, in the size range 250-500 pm, could be controlled to better than 25 pm using the microscope fab- rication method.Furthermore, the technique allows low micrometre control over the gap dimension with estimations of the average gap size along the device length within &1 pm.' The dual-cylinder microelectrodes used in this work consist of a pair of platinum cylindrical electrodes of 25 pm diameter with a microscopic interelectrode spacing of vari- able size (<188 pm).Only devices showing regular geometric features in addition to displaying identical cyclic voltam- metric and chronoamperometric behaviour for the individual cylinder electrodes (i.e. variations in the quasi-steady-state currents ~3%)were used in collection and feedback experi- ments.Table 1 lists the gap dimensions for a series of dual-cylinder microelectrodes under consideration here. Amperometric Measurements Chronoamperometry experiments were carried out using a bipotentiostat, RDE4, Pine Instruments, USA. The generator J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 607 Table 1 Steady-state currents at the generator and collector elec- +-Itrodes for a series of dual-cyclinder devices with variable w 1.0 0.8 -/e3.5 520.0 430.5 45 1.6 15.1 4.7 !! 5.0 469.2 328.2 380.7 23.2 13.8 12.0 339.8 174.3 250.6 35.6 30.4 22.5 250.4 129.2 188.0 33.2 31.3 c35.0 233.7 90.7 156.0 49.8 41.9 , 79.5 194.4 41.1 111.1 75.0 63.0 ii181.5 212.4 32.7 85.1 149.6 61.6 N188.0 185.8 10.4 82.8 124.4 87.4 Experimental data include background subtraction and are normal-ised in cylinder length (L = 1.0 mm) and concentration (1.0 mmol dm-3).Theoretical currents have been calculated from eqn. (2) using experimental values for device dimensions, with D s-', Cb= 1.0 mmol dmP3 and n = 1 for the hexacyanoferrate(III/n) = 7.0 x 10-lo m2 0.2;+ anions in aqueous KCl solution. and collector current-time curves were recorded on an x-y-t recorder 3086 from Hokushin Electric, Japan. Bipotentiosta- tic chronoamperometry of K,Fe(CN), (1.0 mmol dm- 3, was performed in aqueous KCl (0.5 mol dm-3) solution by apply- ing a potential step of +0.4 to 0.0 V us. SCE at the generator electrode to produce Fe(CN):-. The collector electrode was controlled at a constant potential of +0.4 V in order to oxidise hexacyanoferrate(rr) back to hexacyanoferrate(rr1).Three i-t curves were recorded in succession for each experi- ment. The average steady-state current values sample at 10 s are presented in Table 1. The electrochemical cell housed a reference electrode (SCE) and a counter electrode (Pt wire 1 mm in diameter and 10 mm in length) with additional cap fittings to ease handling of the dual-cylinder strip devices. Furthermore, the cell allowed electrolyte deoxygenation, which is often essential for low-concentration experiments and had a maximum solution capacity of 20 ml. The tem- perature in the chronoamperometric experiments reported here was controlled to 26 f1 "C.Results and Discussion Chronoamperometric measurements for the platinum dual- cylinder microelectrodes using the hexacyanoferrate(rrr/Ir) couple are presented in Table 1, where the generator and col- lector steady-state currents are given as a function of the gap dimension. The diffusion currents observed at the generator electrode, i, , are noticeably larger than the corresponding collector electrode currents, i,, for all values of the gap dimension under review. This is an obvious consequence of the generator being supplied with hexacyanoferrate(Ir1) from the electrolyte solution while the collector response is depen- dent on the cross-diffusion of the generator species [hexa- cyanoferrate(rr)]. Furthermore, the experimental values for the generator and collector currents are enhanced and con- verge on diminution of this interelectrode distance, w.The i-w behaviour in the steady-state response is further empha- sised in the recycling efficiency curve shown in Fig. 2 where the recycling efficiency, N, is plotted against the spacing parameter, w, for a series of dual-cylinder microelectrodes. The recycling curve, in addition to illustrating the sensitivity of the iJi, ratio to the gap, indicates that the stationary field recycling model presented earlier,14 where N is unity for any value of w, is not valid over the device dimensions used in this study. The predictions outlined in the stationary-field model, where comparison was made with a microcylinder response, suggested that eqn.(2) might only be valid for a limited case. Nevertheless, the experimental data, which - 1 a ~ I 1 I I I 0.0 40.0 80.0 120.0 160.0 200.0 w/Pm Fig. 2 Recycling curve for a series of dual-cylinder microelectrodes. The plot shows the response of the recycling efficiency, i&, as a function of the micrometer gap spacing, w. demonstrate N values for dual cylinders ranging from 0.06 to 0.83, do show that as the interelectrode distance is reduced (w +0) the device's steady-state response tends to the theo- retical limit, where i, = i,. The i-w data listed in Table 1 also provide a comparison of experimental diffusion currents with their theoretical steady- state values. The result illustrates how well the analytical model predicts the behaviour of this feedback de~ice.~,'~ Strong deviations from eqn.(2) are found in the generator and collector currents for devices possessing large values of the gap spacing, i.e. >5 pm. The current at the generator is also seen to overestimate consistently the values given by the recycling model. However, such deviations become less pro- nounced as the interelectrode gap decreases to a few pm. In contrast, values for the collector current remain below that of the theoretical current since the amperometric response of this electrode is driven by the generator flux with its associ- ated bulk exchanges. For dual-cylinder devices with broad w spacings the collector currents assume low values, typically < 10 nA per mm of cylinder as w approaches 200 pm. For the smallest gap dimension fabricated using the micro-manipulation technique, ca.3.5 pm, the experimental data for i, and i, deviates from the expected values given in eqn. (2) by 15.1% and -4.7%, respectively. Larger deviations follow as the feedback and collection fluxes are minimised over bulk solution fluxes as the collector is placed further from the gen- erator electrode. It is informative to extrapolate this experi- mental data with respect to eqn. (2) in order to understand the performance limitations (e.g. analytical sensitivity, i/C) of dual-cylinder systems. Using i us. l/cosh-'(w) plots it is pos- sible to make estimations of the interelectrode distance neces- sary to approach theory.In this case the generator and collector currents are found to approximate the theoretical currents to better than 5% as the gap dimension diminishes below 2 pm. Furthermore, convergence of experimental cur- rents on the theoretical value is predicted for the present 25 pm cylinder devices as the gap is reduced to CQ. 0.5 pm. Realistically, however, it is unlikely that dual-cylinder devices with such a small dimension could ever be constructed by manual micromanipulation technology. Fig. 3 considers the experimental diffusion currents at dual- and single-cylinder microelectrodes. l4 The steady-state 608 1000.0 800.0 p 600.0 '=-2 t 3 400.0 200.0 .-\r, wc .. I 1, : lo i I I I I I I I I I 0.0 40.0 80.0 120.0 160.0 200.0 Fig.3 The behaviour of the steady-state currents at generator and collector electrodes with varying cylinder gap. The horizontal solid line gives the quasi-stady-state current at a single-cylinder microelec-trode sampled after 10 s pulse. response of the generator electrode is in fact a combination of a cylindrical flux originating from bulk solution and the col-lector feedback process. For small gap sizes (<5 pm), the generator current behaviour is strongly influenced by feed-back effects from the collector electrode which is essentially limiting the spatial extension of the generator diffusion field. Consequently this affords chronoamperometric curves dis-playing a fast onset for steady-state currents in addition to amplification of the diffusion current over that of a single-cylinder device.'0.'3.'4 The generator curve also conveys the effects of the collector electrode positioned at greater distance from the generator.Here, the current at the generator tends to a minimum value which corresponds to single-cylinder behaviour, i.e. i, -,is, as w extends beyond a critical feedback space, wg. At large interelectrode distances therefore the gen-erator response remains relatively unperturbed by electro-chemical processes at the collector electrode. By comparison with the current at a single-cylinder electrode, this distance, wg, is of the order of 80 pm and marks the diffusional dis-tance over which the feedback effect operates for dual-cylinder microelectrodes of this dimension.The collector curve featured in Fig. 3 displays the typical analyte collection properties of dual-cylinder microelectrodes. The current at the collector tends to a background value, in this instance, i, = 0.85 nA, as the interelectrode spacing exceeds ca. 110 pm. Hence, the collector curve defines the maximum collection or sensing distance, w, , beyond which generated species can perturb the collector electrode current from a background value. In these experiments with hexacyanoferrate(rI1) in aqueous KCl solution, the Fe(CN):-species can be detected J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 at the collector over distances > 180 pm from the generator. Further to this, the collector current approaches the theoreti-cal limit with decreasing values of the gap distance.Conclusion Dual-cylinder microelectrodes have been fabricated by micro-scope manipulation procedures to tolerances acceptable for electrochemical characterisation. The amperometric response of this novel feedback device was assessed in terms of the steady-state currents at individual generator and collector microelectrodes. The theoretical treatment previously re-ported set out to explain the diffusion current at dual cylin-ders in terms of a total recycling model. However it was shown here that the model only approximates device per-formance where the interelectrode gap dimension is small relative to the microcylinder diameter. This would be a direct consequence of a high recycling flux between the cylinder electrodes overcoming the effects of bulk solution concentra-tion gradients.The authors acknowledge the co-operation and financial con-tribution of the National Science Foundation of China to the dual-microelectrode projects. B.J.S. wishes to express his gratitude to the students, technical staff and colleagues at Wuhan University for making the visit to China in 1990 rewarding and memorable. References 1 M. S. Harrington and L. B. Anderson, Anal. Chem., 1990, 62, 546. 2 B. J. Seddon, H. H. Girault and M. J. Eddowes, J. Electroanal. Chem., 1989,266,227. 3 K. Aoki, M. Morita, 0. Niwa and H. Tabei, J. Electroanal. Chem., 1988,256,269. 4 H. Tabei, T. Horiuchi, 0. Niwa and M. Morita, J. Electroanal. Chem., 1992,326,339. 5 W. Peng, Z. Zhou and P. Li, J. Electroanal. Chem., 1993,347, 1. 6 W. Peng and B. J. Seddon, Proceedings of the 4th. National Con-ference of Electroanalytical Chemistry, Shanghai, 1990. 7 T. Horiuchi, 0. Niwa, M. Morita and H. Tabei, Anal. Chem., 1992,64,3206. 8 V. T. Shea and A. J. Bard, Anal. Chem., 1987,59,2101. 9 0.Niwa, M. Morita and H. Tabei, Anal. Chem., 1990,62,447. 10 B. J. Seddon, Ph.D. Thesis, Edinburgh University, 1989. 11 W. Peng, Z. Zhou, X. Zhang, B. J. Seddon and Z. Zhao, Fenxi Huaxue, 1992,20,844. 12 Z. Zhou, W. Peng, X. Zhang and B. J. Seddon, Proceedings of the International Beijing Conference and Exhibition on Instrumen-tal Analysis, Science Press, Beijing, 1991, p. F63. 13 C. Wang, B. J. Seddon, W. Peng and X. Zhang, Electrochim. Acta, to be submitted. 14 B. J. Seddon, H. H. Girault, M. J. Eddowes, W. Peng and Z. Zhao, J. Chem. SOC.,Faraday Trans., 1991,87,2603. Paper 3/06341K; Received 25th October, 1993

ISSN:0956-5000    

DOI:10.1039/FT9949000605

出版商:RSC    

年代:1994

数据来源: RSC

摘要:

J. CHEM. SOC. FARADAY TRANS., 1994, 90(4), 609-615 Voltammetric and Subtractively Normalized Interfacial FTlR Study of the Adsorption and Oxidation of L( +)=Ascorbic Acid on Pt Electrodes in Acid Medium: Effect of Bi Adatoms Miguel A. Climent,t Antonio Rodes, Maria J. Valls, Juan M. Perez, Juan M. Feliu and Antonio Aldaz Departamento de Quimica Fisica , Universidad de Alicante, Apartado 99,03080 Alicante, Spain The adsorption and oxidation of L(+)-ascorbic acid on bare and Bi-covered Pt electrodes in a sulfuric acid medium have been studied by cyclic voltammetry and in situ FTlR spectroscopy. The latter technique has allowed the direct detection of CO,,, from ascorbic acid adsorption, both at opencircuit and controlled potential, for the bare Pt surface. Under these latter conditions both linearly bonded and multi-bonded CO were detected.The existence of other strongly adsorbed species which are oxidized in the oxygen adsorption region has been demonstrated by cyclic voltammetry. The formation of all these adsorbed species is completely suppressed on the Bi-covered Pt surface under UPD conditions. Ascorbic acid oxidation was monitored by following the varia- tion of the IR spectra in the presence and in the absence of Bi3+ in solution. The electrocatalytic effect of the 6i adlayer can be explained mainly through the suppression of any dissociative adsorption step leading to the formation of strongly adsorbed species. Vitamin C, L(+)-ascorbic acid (AA), is a compound that takes part in many important life processes.Some of its bio- chemical functions, although they are not yet known precisely, are clearly related to its reductor character: AA is very easily oxidized, chemically or electrochemically, to dehydro-L( +)-ascorbic acid (DHA). This latter species undergoes a hydra- tion reaction leading to the hydrated product (HDHA). The electrochemical oxidation of AA on platinum has been studied less than that on or gold7p8 electrodes. The reversible, non-destructive, adsorption of AA at a Pt surface, partially blocking the hydrogen adsorption sites, was first reported by Brezina et aL9 These researchers suggested that this adsorption is responsible for the high overvoltage of the AA oxidation reaction on Pt, that proceeds in acid media via a (2H+, 2e-) irreversible process giving DHA; the adsorbed species would be oxidized in the potential range of surface oxide formation.Aldaz et a1." pointed out the forma- tion of a radical intermediate during the electrochemical oxi- dation of ascorbate on Pt and other electrodes. Karabinas et a1.,l1 on the basis of kinetic parameters, proposed a mecha- nism for the electro-oxidation of AA on Pt, suggesting the presence of anionic adsorbed intermediates ;the desorption of the product (DHA) would be the rate-determining step. Recently," the structure-sensitive character of AA electro-chemical oxidation has been demonstrated by using Pt single-crystal electrodes. Electrochemical evidence has also been given for the existence of spontaneous dissociative adsorption of AA on a Pt(100) surface.Under these condi- tions a strongly adsorbed species is formed, that is oxida- tively stripped at the same potential as CO (0.78 V).12 Therefore the oxidation of AA on Pt electrodes would consist of two pathways: direct oxidation to DHA without adsorp- tion, and another mechanism leading to CO, as a final product through dissociative adsorption of AA and forma- tion of CO as an adsorbed intermediate.', Formation of the same species on a polycrystalline Pt electrode has been con- firmed spectroscopically in an in situ Fourier-transform infra- red reflection-absorption spectroscopy (FTIRRAS) study by Xing et al.' The catalytic effects of several metal ions on the oxidation of AA on Pt electrodes were first reported by Takamura and t Permanent address: Department de Ingeniena de la Con-struccion, Obras Publicas e Infraestructura Urbana, Universidad de Alicante, Alicante, Spain.Sakam0t0.l~ These metal ions were adsorbed in under-potential deposition (UPD) conditions. It has been shown that single-crystal Pt(h, k, r) surfaces modified by irreversibly adsorbed Bi catalyse the oxidation of AA, changing the mechanism to a structure-insensitive process.12 This effect was attributed to the suppression of the dissociative adsorp- tion pathway.', The oxidation of AA on gold electrodes is also a structure-insensitive process and does not produce adsorbed CO.lS The aim of this work is to study the adsorption and oxida- tion of AA on polyoriented Pt bare electrodes and on the corresponding Bi-covered Pt surfaces, by cyclic voltammetry and SNIFTIRS (subtractively normalized interfacial Fourier- transform infrared spectroscopy).The study will be focused on the effect of the dissociative adsorption pathway on the overall AA electrooxidation mechanism. Experimental Solutions were prepared by dissolving the appropriate amount of L(+)-ascorbic acid (Merck p.a.) in 0.1 mol dm-3 H,SO, (Merck suprapur) as test electrolyte. Ultrapure water was supplied by a Millipore Milli-Q system. Deaeration of solutions was accomplished by argon bubbling. Spherical single-crystal platinum electrodes (polyoriented surface) were used for electrochemical experiments. Before each experiment the platinum samples were cleaned by heating in a gas-oxygen flame.They were then cooled in air, protected with a water droplet in equilibrium with air and transferred to the ~11.'~Other experiments were performed with a polycrystalline electrode extensively cycled between 0.06 and 1.4 V to compare these results with those obtained in the spectroelectrochemical cell, in which a polycrystalline electrode was employed. The potentials were measured against a reversible hydrogen electrode (RHE). Open-circuit AA adsorption experiments were performed following the experimental procedure described by Clavilier and Sun," without air exposure of the adsorbed species, as described below. The Bi-modified Pt electrodes were obtained by irreversible adsorption from a Bi3+ solution prepared by dissolving Bi,O, (Merck p.a.) in 0.1 mol dm-3 H2S0,.The clean elec- trodes were immersed in the Bi3+-containing solution for 30-120 s rinsed with water and immersed in the test solution, free of Bi3+ ions. Thus the adatoms were present only on the electrode surface. Alternatively, experiments were carried out in Bi3+-containing solutions, in a classical UPD pro- cedure.l4 The spectrometer employed was a Nicolet 5PC instrument with a liquid-nitrogen-cooled MCT detector. A glass external-reflection cell with a CaF, window was used for the spectroelectrochemical experiments. The bulk polycrystalline platinum working electrode mounted on a glass rod was pol- ished mechanically with successively smaller grades (1 .O, 0.3 and 0.05 pm) of alumina, and further steam washed.This electrode was then cycled between 0.06 and 1.6 V us. RHE until a stationary blank voltammogram was obtained. A platinum counter electrode was used; a saturated mercury(1) sulfate electrode was used as reference, but all measured potentials were converted to the RHE scale. The cell com- partment of the spectrometer was purged throughout the experiment by flowing clean air, free of CO, and H,O. The incident angle at the electrode surface was close to 60" with respect to the surface normal. Two types of FTIR spectra were obtained: (a) absolute spectra in which the reference spectrum was recorded with a 0.1 mol dm-3 H2S04 solution as test electrolyte.The sample spectrum was taken after adding the appropriate amount of AA. In both cases open-circuit conditions were employed. (b) Potential difference spectra were obtained by recording refer- ence and sample at different potentials with the same AA solution. Typically 920 spectra (approximately 2 scans s-', 8 cm-' resolution), were recorded at each potential with the voltage being switched every 92 scans. In order to monitor the AA oxidation process FTIR mea- surements were taken during slow potential sweeps (1 mV s-') controlled by software. This allows the acquisition of 99 interferograms in 50 s, corresponding to a potential window of 50 mV. The IR spectra obtained from these interferograms were ascribed to the centre of the potential interval where they were recorded.The first spectrum of the series was taken as the reference for the whole set. Results and Discussion Electrochemical Experiments Fig. l(a) shows the voltammetric behaviour of a spherical single-crystal Pt electrode in contact with a 0.1 mol dm-3 H2S04 solution during the first potential cycle including adsorption of an oxygen monolayer. The voltammograms corresponding to the first oxidation cycle of AA on the same f3 Fig. 1 Voltammograms of a spherical single-crystal Pt electrode in 0.1 mol dm-3 H,SO, + X mol dm-' AA solutions. X =(a),0, (b) 5.6 x (c) 4.7 x Sweep rate, I) = 50 mV s-'. J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 electrode, at 5.6 x lop4 and 4.7 x lo-, mol dm-3 concen- trations, are also shown in Fig.l(b) and (c). Before these volt- ammetric curves were obtained the flame-treated electrode was placed in contact with the solution at 0.5 V. This poten- tial was held for 30 s in order to allow adsorption of AA from solution and then the potential was cycled between 0.06 and 0.5 V to determine the degree of suppression of hydrogen adsorption &,, (Sb = 1 -S,3, due to AA adsorption. The Sb values obtained were 0.57 and 0.66 for the low and high AA concentrations, respectively, in good agreement with pre- viously reported result^.^ It may be appreciated from Fig. 1 that AA adsorption takes place approximately to the same extent on all crystallographic orientations present on a spher- ical single-crystal Pt electrode.On a spherical single-crystal Pt electrode AA oxidation starts at potentials slightly higher than 0.5 V, i.e. before surface oxide formation [Fig. l(b)]. An additional surface process is observed at 0.79 V identical to that reported for the case of AA oxidation on Pt(100).12 The main AA oxidation peak on a spherical single-crystal Pt electrode appears at 0.9 and 1.0 V for the 5.6 x and 4.7 x lo-, mol dm-3 AA concentrations. Note, that in Fig. l(c), the higher oxidation density current is observed in the negative-going sweep rather than the positive-going sweep, at potentials below 0.8 V. This behaviour was also observed with Pt( 11 1) and Pt( 110) electrodes.12 This fact suggests the releasing of active sites upon oxidation of some adsorbed inhibiting species at high potentials.In order to isolate the irreversibly adsorbed species formed on the surface of the spherical single-crystal Pt electrode, open-circuit AA adsorption experiments l7 were performed. In Fig. 2 the various steps required for the electrochemical characterization of adsorbed species are reported : (a) the hydrogen adsorption-desorption profile of the bare Pt elec-trode in contact with the test electrolyte. (b) The hydrogen adsorption-desorption profile of the surface saturated with the adsorbed species, with 0.06 and 0.5 V as potential limits. This curve was recorded after putting the electrode in contact, for 1 min, with a 0.1 mol dm-3 AA-0.1 mol dm-3 H2S04 solution and then transferring to the test electrolyte without air exposure.(c) The oxidative stripping of the adsorbed species which leads to voltammetric profile with two marked peaks at 0.735 and 0.795 V. (d) The recovery of the hydrogen adsorption capability after oxidation of the adsorbed species. A comparison of the voltammograms shown in Fig. 2 with those reported for the dissociative adsorption of formic acid and methanol on platinum electrode^,'^ strongly suggests that the same irreversibly adsorbed species are formed in the adsorption processes of AA and of the smaller organic mol- ecules. From the potential range where this oxidative strip- ping takes place [see Fig. 2(c)], the adsorbed species can be identified as CO. The existence of two oxidation peaks at 0.735 and 0.795 V can be related to the structural heter- ogeneity of the electrode surface.As stated for formic acid and methan01,'~ the first peak can be ascribed to the strip- ping of the species adsorbed on (111) or (110) sites and the second to the same process on (100) sites. Quantitative analysis of the voltammetric results of Fig. 2 has been performed following ref. 17. The percentage of hydrogen sites blocked by the adsorbed species, S,, the per- centage of hydrogen sites recovered after its oxidative strip- ping, s,,and the number of electrons, n, transferred per Pt surface site during the latter process are 62%, 84% and 1.67 respectively. The n value obtained is consistent with the for- mation of both linearly bonded and bridge-bonded CO on the electrode surface.However the rather low value obtained for S,, in contrast with the 96% found for Pt(100)," suggests J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 ! I, I'80 I' NI E 4O > E/V vs. RHE -80 Fig. 2 Voltammograms obtained with a spherical single-crystal Pt electrode in the different steps of a dissociative adsorption experi-ment with AA. (a) Voltammogram of bare electrode. (b) Hydrogen adsorption-desorption profile of the surface saturated with the che-misorbed species. (c) Oxidation of the chemisorbed species. (d) Hydrogen adsortion-desorption after oxidation of the chemisorbed species. Test electrolyte: 0.1 mol dm-3 H,SO,. u = 50 mV s-'. Immersion potential 0.5 V.that other species are adsorbed on the polyoriented Pt surface.These species need higher potentials for oxidation. Another open-circuit adsorption and isolation experiment was performed with a polycrystalline Pt electrode extensively cycled between 0.06 and 1.4 V. Under these conditions no additional changes of the electrode surface structure are expected during the stripping of the adsorbed species which oxidizes in the platinum oxide region." At the same time the surface structure conditions achieved in this way are similar to those reached in the FTIR experiments. Fig. 3 shows the results obtained. It may be appreciated from curve (c) that oxidation of two types of isolated adsorbed species occurs: 611 CO-like species that are stripped from the surface in a poten-tial range 0.6-0.9 V and other species that are oxidized at potentials higher than 0.95 V.The voltammetric curve corre-sponding to oxidation of adsorbed CO on this perturbed electrode is less resolved than the one corresponding to the polyoriented Pt surface (Fig. 2), showing the effect of the breaking of long-range ordered domains after extensive elec-trochemical cycling. The comparison of curves (a) and (d)in Fig. 2 and 3 shows clearly that a higher S, value is obtained in the latter case. The oxidation at potentials where the Pt surface is covered by oxygen is typical behaviour of unsaturated molecules that are adsorbed on Pt and, in principle, are able to form surface n-complexes" (e.g.ethene,22acetylenez3and acetonez4).The electrochemical results reported here do not allow determi-nation of the nature and origin of the adsorbed species that are oxidized in the Pt oxide region. They may be formed in the same dissociative adsorption process of AA leading to COads,or they may be the result of a different interaction between the AA molecule and the electrode surface (e.g. a different spatial orientation). However, the effect of the surface anisotropy in promoting the formation of such adsorbed species is clear. This can be seen by comparing the S, values with Pt(100) and polyoriented Pt after stripping the adsorbed species that are oxidized below 0.9 V (see preceding paragraphs). Ascorbic Acid Oxidation on a Bi-couered Pt Electrode The voltammograms shown in Fig.4 were obtained with a spherical single-crystal Pt electrode modified by irreversibly adsorbed Bi at high coverage, OBi = 0.86, (OBj = 1 -OH). Curve (a)corresponds to the Bi-modified electrode in contact with the test electrolyte, in the absence of AA. In this curve, the voltammetric peaks appearing at 0.63 and 0.9 V corre-spond to surface processes related to the presence of the Bi adlayer.' * Those at the higher potentials involve a partial desorption of the adlayer with a releasing of hydrogen adsorption sites. AA oxidation on a Bi-modified (irreversibly adsorbed) spherical single-crystal Pt electrode produces a single wave with peak potential at 0.615 and 0.68 V for the 8.1 x and 6.8 x mol dm-3 AA concentrations, respectively, Fig.4(b) and (c). These peak potentials are in good agreement with those obtained for AA oxidation on the three Pt basal planes highly covered with Bi. This confirms the structure-(C) --* i 2 mA cm-2 ju -100 t Fig. 4 Voltammograms of a spherical single-crystal Pt electrode Fig. 3 Voltammograms corresponding to the dissociative adsorp-highly covered with irreversibly adsorbed Bi (OBi = 0.86), in 0.1 mol tion of AA on a polycrystalline Pt electrode extensively cycled dm-3 H,SO, + X mol dm-3AA solutions. X = (a) 0,(b)8.1 x loe4 between 0.06 and 1.4 V. Same conditions as in Fig. 2. (c) 6.4 x lo-'. u = 50 mV s-'. insensitive character of the AA oxidation under these condi- tions.12 A comparison of curves (b) and (c) in Fig.1 and 4 shows that the AA electrocatalysis produced by the Bi adlayer is substantiated in a 0.3 V decrease of the peak potentials and a slight increase of the peak current densities for the Bi- modified electrodes. Nevertheless, it can be appreciated that the electrocatalytic effect of Bi is less effective in the potential region of the foot of the AA oxidation wave: this electro- chemical reaction starts at a potential only slightly higher (0.05-0.1 V) on the Pt bare electrodes than on the corre- sponding Bi-modified Pt surface. Note that the AA oxidation wave in Fig. 4(c) is practically identical to that obtained with a polycrystalline Pt electrode in contact with a Bi3 +-containing solution. This indicates that the electrocatalytic effect of Bi adatoms toward AA oxi-dation is independent of the way the adlayer is obtained: irre- versible adsorption or UPD from solution, provided that a high Bi coverage is reached on the Pt surface. Open-circuit AA adsorption experiments have also been performed on Bi-modified Pt electrodes.Fig. 5, solid line, shows the first and second voltammetric cycles obtained in the test electrolyte with a spherical single-crystal Pt electrode highly covered with irreversibly adsorbed Bi. The upper potential limits were 0.75 and 1.4 V, respectively. The broken line in the same figure shows the behaviour of the same elec- trode newly covered with Bi, and put in contact with an AA solution (same conditions as in the experiments reported in Fig.2 and 3). It can be deduced from the electrochemical behaviour shown in Fig. 5 that no irreversibly adsorbed species have been formed on the Bi-modified Pt surface since both experiments led to the same voltammetric curves. The small differences in current density can be attributed to the slightly different Bi coverages reached in each experiment. In this way it can be stated that AA does not undergo sponta- neous dissociative adsorption, and consequently CO is not produced on the Bi-modified Pt-polyoriented electrode. Spectroelectrochemical Experiments Fig. 6 shows the in situ FTIR absolute spectrum obtained with the polycrystalline Pt electrode in contact with a 0.05 rnol dm-3 AA-0.1 mol dm-3 H2S04 solution at open-circuit, with the incident beam in the p polarization plane.This spectrum represents the difference between those obtained with the aforementioned solution and with the test electrolyte. It is interesting to note that the spectrum shown in Fig. 6 has been collected under the same conditions, i.e. at open-circuit, as those used for the isolation of the strongly a 'r E/V vs. RHE -1 00 -t Fig. 5 Solid line: voltammogram of a spherical single-crystal Pt electrode highly covered with irreversibly adsorbed Bi (near monolayer) in 0.1 rnol dm-3 H,SO,. Broken line: similar experi- ment with a newly Bi-covered surface, but after contacting the elec- trode with an AA + H,SO, solution. D = 50 mV s-'. Immersion potential: 0.1 V. (See text for other conditions.) J.CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 2200 2000 1800 1600 1400 1200 1000 waven u m ber/cm -' Fig. 6 Absolute FTIR spectrum obtained with a polycrystalline Pt electrode in contact with a 0.1 rnol dm-3 H,SO, + 0.05 rnol dm-3 AA solution. Reference spectrum recorded with the test electrolyte. Open-circuit conditions, p polarization. adsorbed intermediates in the experiments reported in Fig. 2 and 3. In Fig. 6 a band appears, centred at 2050 cm-', that corresponds to on-top adsorbed CO. As expected this band is not observed when p-polarized is substituted by s-polarized radiation (spectrum not shown). The remaining bands, that are common to the spectra obtained with p and s polariza-tions are due to species in solution.The strong band centred at 1765 cm-' can be assigned to the C-0 stretching vibra- tion of AA in sol~tion,~~~~~ while that at around 1700 cm-' can be originated from C=O and/or C-C stretching contri- butions of AA or related anions.25 The strongly coupled 0-H bending and the CH, deformation modes of AA, produce the very broad band observed between 1500 and 1250 cm-'. Finally, the bands below 1250 cm-'are assigned to C-0 and C-C stretching modes.26 Fig. 7 shows potential difference FTIR spectra obtained with a polycrystalline Pt electrode in contact with a 0.05 mol dm-j AA-0.1 rnol dm-3 H2S04 solution. Spectrum (a)was obtained without Bi3+ in solution, while spectra (b)and (c) were recorded in the presence of 1 x and 2 x mol dm-3 Bi3+ , respectively.In these spectra reference and sample potentials are within the potential region where the I 2200 2100 2000 1900 1800 wavenumber/cm -' Fig. 7 Potential difference spectra obtained with a polycrystalline Pt electrode in contact with a 0.1 mol dm-j H,SO, + 0.05 mol dm-3 AA solution + X rnol dm-3 Bi3+ solution. X =(a) 0, (b) 1 x lo-,, (c) 2 x Sample potential: 0.435 V, reference poten- tial: 0.055 V US. RHE. J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 bare platinum electrode is blocked by the adsorbed species coming from AA (see Fig. 1). On the other hand, in this potential range, the electrode surface is expected to be covered with a Bi adlayer in the Bi3+ containing solution (see Fig. 4 and 5). Spectrum (a) shows a bipolar band, centred at around 2050 cm-', characteristic of linearly adsorbed CO in acid medium.27 This feature was observed for AA concentra-tions down to 3 x mol dm-3.As expected, the decrease of AA concentration produces a lower CO coverage of the electrode, giving rise to a signal shifting toward lower wave- number^,^^.^^ and a decrease in the CO band intensity. It is interesting to note in Fig. 7(a) the presence of a band at ca. 1850 cm-', characteristic of multi-bonded CO. As can be observed in Fig. 7(b) and (c), the presence of Bi3+ ions in solution inhibits the formation of CO on the Pt surface. When the Bi3+ concentration reaches 2 x mol dm-3 the CO signal cannot be detected. This result is in agreement with the voltammetric experiments shown in Fig.4 and 5. Fig. 8 shows the voltammetric curves obtained in the spec- troelectrochemical cell: (a)is the blank voltammogram, (b)is the voltammetric curve in presence of 0.05 mol dmm3 AA and (c) corresponds to the same AA solution to which 2 x mol dm -Bi3+ has been added. Both curves (b)and (c)were obtained with the Pt electrode pushed against the CaF, window, just after collecting spectra (a) and (c) in Fig. 7. These AA oxidation voltammetric profiles are comparable to those shown in Fig. l(c) and qc), taking into account the different mass-transfer conditions existing in a thin-layer spectroelectrochemical cell (Fig. 8) and in an ordinary elec- trochemical cell (Fig. 1 and 4). , I' N a t Fig.8 Voltammograms obtained with the polycrystalline Pt elec- trode in the spectroelectrochemical cell. (a) 0.1 rnol dm-3 H,SO, solution, u = 100 mV s-l, electrode separated from the CaF, window. (b) 0.1 mol dm-3 H,SO, + 0.05 rnol dm-j AA solution, u = 1 mV s-', electrode pressed against window. (c) Same conditions and solution of (b) but with the addition of 2 x rnol dm-3 Bi3+. 613 Fig. 9(a) shows a set of potential difference spectra obtained with the polycrystalline Pt electrode during the potential sweep shown in Fig. 8(b).The reference was the first spectrum of the series that corresponded to a potential of 0.06 V. The spectrum corresponding to 0.31 V shows clearly the bipolar band centred at 2050 cm-', characteristic of lin- early bonded CO in acid medium.The 0.51 and 0.56 V spectra, recorded in the potential region at the foot of the AA oxidation voltammetric wave, show the simultaneous appear- ance of two negative-going absorption bands around 1770 and 1700 cm-', respectively, and a positive-going band at 1800 cm-'. The negative-going bands are related to the dis- appearance of the electrochemical oxidation reactant (AA), while the positive-going one can be attributed to the C=O stretching of the final oxidation product, in its hydrated bicyclic form (HDHA), as indicated by Xing et al.' Another negative-going absorption band can be seen near 1400 an-' due to AA consumption. Several positive-going bands appear below 1300 cm-'. This zone of the spectrum, which contains the strong absorption band of bisulfate anions at 1200 cm-', is highly mixed, thus preventing a clear assignment.The oxidation of adsorbed CO takes place at potentials positive to the beginning of AA oxidation. This can be moni- tored by the decrease of intensity of the 2050 cm-' bipolar band and by the appearance of the CO, absorption band at 2345 cm-'. The latter is first observed in the 0.66 V spec-trum, [Fig. 9(a)], in agreement with the CO oxidation volt- ammetric experiment shown in Fig. 2(c). Fig. 1qa) represents a plot of the integrated intensity of the CO, band (2345 cm-') in the spectra obtained in the poten- tial sweep in Fig. 8(b). It is observed that CO, production starts at approximately 0.6 V. The CO, band intensity remains almost constant in the potential range between 0.7 and 1.0 V approximately, and increases for potentials higher than 1.0 V.This evidences a change either in the kinetics or in the mechanism of the C0,-forming pathways. CO, can be produced from oxidation of the adsorbed CO and from further oxidation of the AA oxidation product^'^ and resi- dues of the AA dissociative adsorption on the Pt surface. No spectroscopic evidence of the presence and/or nature of the adsorbed species that oxidize in the Pt oxide potential region, (Fig. 3), can be drawn from the spectra obtained in this work. Fig. 9(b) shows a set of potential difference spectra similar to that reported in Fig. 9(a) but in a Bi3+-containing AA solution. These spectra were collected during the potential sweep shown in Fig.8(c). In the same way, Fig. 1qb) rep-resents the plot of the integrated intensity of the CO, band in this spectra series. All bands in the spectra in Fig. 9(a) are qualitatively similar, except obviously the CO signal, that is absent in the spectra corresponding to the Bi-covered poly- crystalline Pt (see Fig. 7). A comparison of the spectra corre- sponding to potentials up to 0.81 V allows spectroscopic confirmation of the following points: (a) AA oxidation starts at potentials slightly lower on the Bi-covered Pt electrode than on the bare Pt surface. (b) The direct AA oxidation pathway on both electrodes produces neither CO nor C02 in the potential range 0.4-0.83 V. The only C0,-forming pathway that works in the potential range mentioned above is the oxidation of CO formed by AA dissociative adsorption.The CO, production in the AA oxidation mechanism on the Bi-covered polycrystalline Pt electrode starts at approx- imately 0.86 V, as can be seen in Fig. 9(b) and lqb). For potentials higher than 1.1 V the CO, absorption band is more intense in the Fig. 9(b) spectra than in the Fig. 9(a) spectra, see Fig. 10. The Bi adlayer starts to be desorbed elec- trochemically at potentials around 0.9 V [see Fig. 4(a) and 5); so the CO, formation via AA dissociative adsorption is J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 I 2500 2000 1500 1000 2500 2000 1500 1000 wavenumber/cm-' wavenumber/cm-' Fig. 9 (a) Set of SNIFTIR spectra collected during the potential sweep of Fig.8(b).Reference spectrum corresponds to 0.06 V. See text for details. (6) Spectra taken during the potential sweep of Fig. 8(c).The numbers on the spectra correspond to E/V us. RHE. 1.2q5 1.0 -0.8-> 4-.-ln d 0.6-.-0.4-0.2 -0 0 0.2 0.4 0.6 0.8 1.0 1.2 1.4 1.6 E/V vs. RHE Fig. 10 Plots of the integrated intensity of the CO, band (arbitrary units) in the spectra series recorded during the potential sweeps of Fig. 8. (a) AA solution without Bi [corresponds to potential sweep of Fig. 8(b) and to 9(a) spectra series]. (b) AA-Bi3+ solution [corresponds to potential sweep of Fig. 8(c) and to 9(b) spectra series]. partially opened at this potential for the uncovered regions of the Pt electrode.Nevertheless, the higher CO, production rate observed at potentials higher than 1.0 V in Fig. lqb), as compared with that in curve (a), suggests the existence of another C0,-forming pathway that performs better in the presence of bismuth. In this way Takamura and Saka-mot~,'~~in an NMR analysis of the AA oxidation products on polycrystalline Pt with and without Bi3+ ions in solution found, in the former case, some new signals that were ascribed to degradation products of DHA : 2,3-diketo-~-gulonic acid (DKG) and its enol form (KGA); these products would be oxidized to CO,. So it may be postulated that the presence of Bi adatoms on Pt or Bi3+ ions in solution promote the further degradation of primary products of the AA oxidation. Conclusions The results obtained in this work confirm that the high over- voltage of AA oxidation on polycrystalline Pt electrodes is due to the presence of irreversibly adsorbed species: linearly bonded and multi-bonded CO that is formed through AA spontaneous dissociative adsorption and other species, pro- duced also when contacting the Pt surface with an AA acidic solution at open-circuit.These latter species, whose nature and origin is not yet known, oxidize in the Pt oxide region. J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 615 As previously suggested for Pt single-crystal electrodes,’ * the electrocatalytic effect, towards AA oxidation, of the Bi adlayer formed on the polycrystalline Pt surface is due to the suppression of any irreversible adsorption step, i.e.to the ability of the Bi-covered electrode to keep all active sites free 9 10 11 12 M. Brezina, J. Koryta, T. Loucka, D. Marsikova and J. Pradac, J. Electroanal. Chem., 1972,40, 13. A. Aldaz and A. M. Alquie, J. Electroanal. Chem., 1973,47, 532. P. Karabinas and D. Jannakoudakis, J. Electroanal. Chem., 1984,160,159. M. J. Valls, J. M. Feliu, A. Aldaz, M. A. Climent and J. Clavilier, for AA oxidation. Nevertheless, a small electronic effect of Bi J. Electroanal. Chem., 1989, 260,237. cannot be ruled out since AA oxidation starts, on the Bi- covered polycrystalline Pt electrode, at potentials slightly lower than on the corresponding bare Pt surface. A detailed study concerning the effect of Bi coverage on the electro- catalysis of AA oxidation would be necessary to clarify this point.Further work is in progress. 13 14 15 16 X. Xing, I. T. Bae, M. Shao and C. C. Liu, J. Electroanal. Chem., 1993,346,309. (a)K. Takamura and M. Sakamoto, Chem. Pharm. Bull. (Tokyo), 1979, 27, 254; (b) K. Takamura and M. Sakamoto, J. Electro-anal. Chem., 1980, 113, 273. X.Xing, M. Shao, M. W. Hsiao, R. R. Adzic and C. C. Liu, J. Electroanal. Chem., 1992,339, 2 1 1. J. Clavilier, R. Faure, G. Guinet and R. Durand, J. Electroanal. The financial support of the DGICYT through contracts PB90-0560 and CE9 1.O001 is gratefully acknowledged. The authors thank the Conselleria d’Educaci6 i Ciencia de la Generalitat Valenciana for the funds for the FTIR facility 17 18 19 Chem., 1980,107,205.J. Clavilier and S.G. Sun, J. Electroanal. Chem., 1986, 199,471. J. Clavilier, J. M. Feliu and A. Aldaz, J. Electroanal. Chem., 1988,243,419. S. G. Sun, These de Doctorat d’Etat, Paris, 1986. purchase. We are also indebted to Prof. J. Clavilier for some 20 J. Clavilier and D. Armand, J. Electroanal. Chem., 1986, 199, valuable comments. 187. 21 A. Wieckowski, Electrochim. Acta, 1981,26, 1121. 22 W. E. Triaca, T. Rabockai and A. J. Arvia, J. Electrochem. SOC., References 1979,126,218. 1 E. Kodicek and K. Wenig, Nature (London), 1938, 142, 35. 2 D. M. H. Kern, J. Am. Chem. SOC., 1954,76, 1011. 3 S. P. Perone and W. J. Kretlow, Anal. Chem., 1966,38, 1760. 4 J. J. Ruiz, A. Aldaz and M. Dominguez, Can. J. Chem., 1977, 55, 2799. 5 J. J. Ruiz, A. Aldaz and M. Dominguez, Can. J. Chem., 1978,56, 1533. 6 J. J. Ruiz, J. M. Rodriguez-Mellado, M. Dominguez and A. Aldaz, J.Chem. SOC., Faraday Trans. 1, 1989,85, 1567. 7 M. Rueda, A. Aldaz and F. Sanchez-Burgos, Electrochim. Acta, 1978,23,419. 23 24 25 26 27 28 29 A. B. Delgado, A. M. Castro Luna, W. E. Triaca and A. J. Arvia, J. Electrochem. SOC., 1982, 129, 1493. B. Bansch., Th. Hartung, H. Baltruschat and J. Heitbaum, J. Electroanal. Chem., 1989,259, 207. J. Hvoslef and P. Klaeboe, Acta Chem. Scand., 1971,25,3043. H. A. Tajmir-Riahi and D. M. Boghai, J. Znorg. Biochem., 1992, 45, 73. B. Beden, A. Bewick, K. Kunimatsu and C. Lamy, J. Electro-anal. Chem., 1981,121,343; 1982,142,435. M. Schemer, Surf: Sci., 1979, 81, 562. A. V. Myshlyavtsev and V.P. Zhdanov, Langmuir, 1993,9, 1290. 8 C. Acerete, L. Garrigos, J. Guilleme, E. Diez and A. Aldaz, Elec-trochim. Acta, 1981, 26, 1041. Paper 3104941H ;Received 16th August, 1993

ISSN:0956-5000    

DOI:10.1039/FT9949000609

出版商:RSC    

年代:1994

数据来源: RSC