摘要:

J. Chem. SOC., Faraday Trans. I, 1985, 81, 1467-1470 Kinetic Electron Spin Resonance Investigation of the Monohydronitro Free Radical of 2,3,5,6-Tetrachloroni trobenzene BY LESLIE H. SUTCLIEFE School of Chemistry, The University, Liverpool L69 3BX Received 24th Spptember, 1984 The e.s.r. spectrum of the short-lived monohydronitro radical from 2,3,5,6-tetrachloronitro- benzene has been obtained by ultraviolet photolysis. The kinetics and energetics of radical recombination have been measured with the aid of two techniques: (i) switching off the radiation from a 1 kW ultraviolet lamp and (ii) pulses from an ultraviolet laser. The two sets of data are in good agreement. We have shown1v2 that the e.s.r. spectra obtained by ultraviolet irradiation of nitrobenzenes in ethers or alcohols originate from the adduct radicals ArN(0)OS ; these arylalkoxynitroxides are formed by addition of solvent radicals S to the acceptor nitro compound^.^ However, the excited nitrobenzene can abstract hydrogen from the solvent. Asmus et aL4 have reported the ultraviolet absorption spectrum and a second-order decay constant (6 x lo8 dm3 mol-l s-l) for the monohydronitrobenzene radical, obtained by pulse radiolysis of nitrobenzene in aqueous acid.This finding led us5 to look for and discover a low stationary-state concentration of monohydronitro radicals formed by initial hydrogen abstraction from the solvent by electronically excited nitrobenzenes. We found that the radicals can be detected only for 2,4,6- trichloronitrobenzene and 2,3,5,6-tetrachloronitrobenzene. Weak e.s.r.spectra of the radicals were obtained from concentrations estimated to be ca. 5 x lop6 dm3 mol-l, and a second-order decay rate constant was estimated to have an upper limit of lo6 dm3 mol-l s-l. The photolysis of chloronitrobenzenes in solution has been shown2 to be complex, the natural decay rate of ArN(0)OH being an important factor. The present paper reports a detailed kinetic e.s.r. study of the decay rate that was made using an ultraviolet laser flash and also a 1 kW continuous ultraviolet source. EXPERIMENTAL MATERIALS Koch-Light laboratories pure-grade 2,3,5,6-tetrachloronitrobenzene was recrystallised twice from warm aqueous ethanol to give long colourless needles. All solvents used were Fisher Scientific Co. reagent grade; no purification was attempted.E.S.R. SPECTROSCOPY A Varian E-4 X-band spectrometer was used in conjunction with a Varian E 500 gaussmeter and a Systron Donner model 6057 microwave frequency counter. PHOTOLYSIS AND KINETICS Continuous ultraviolet illumination was provided by a Hanovia 1 kW lamp type 977B0010, the output being filtered with a cobalt(Ir)/nickel(rI) solution and a 3 mm thick borosilicate glass filter. E.s.r. parameters were obtained from continuously irradiated solutions of 2,3,5,6- 14671468 E.S.R. STUDY OF THE MONOHYDRONITRO RADICAL Table 1. E.s.r. hyperfine coupling constantsa (mT) of the monohydronitro radical, ArN(O)OH, of 2,3,5,6-tetrachloronitrobenzene in propan-2-01 T/"C UN 18.0 7.0 - 13.0 - 23.5 - 32.0 - 40.2 - 50.4 - 58.8 - 67.6 2.3 10 & 0.016 2.310 k0.014 2.308+0.010 2.3 10 k 0.008 2.308 k 0.01 2 2.302&0.010 2.3 12 k 0.002 2.3 10 k 0.012 2.302 & 0.002 0.470 & 0.002 0.466 f 0.002 0.464 & 0.006 0.462 & 0.002 0.462 & 0.004 0.462 & 0.006 0.462 0.002 0.456 & 0.008 0.440 0.002 a Error limits are one standard deviation.tetrachloronitrobenzene: a shutter in front of the lamp enabled decay rates of the radical to be measured. Kinetic measurements were also made with the aid of a Molectron Corporation pulsed nitrogen laser model UW-24: pulse rates of ca. 0.2 to 5 s-' were used. In order to increase the signal-to-noise ratio, multiple scans ranging from 20 to 200 were stored in a Nicolet 1170 computer and the output plotted on a Hewlett-Packard 7005B x-y chart recorder. Soiutions of the substrate (ca. 3 x dm3 mol-l) were degassed before use with four freeze-pump-thaw cycles on a vacuum system.Radical concentrations were estimated by comparing the intensity of radical e.s.r. signals with those from a standard solution of a,a'-diphenyl-P-picrylhydrazyl (DPPH). RESULTS AND DISCUSSION The availability of improved instrumentation compared with our previous investigation5 led us to attempt to generate the monohydronitro radical from 2,3,5,6-tetrachloronitrobenzene in a variety of solvents. These were toluene, anisole, benzyl methyl ether, 1 -phenylethanol, tetrahydropyran and propan-2-01. Only the latter solvent produced the monohydronitro radical but tetrahydropyran gave an e.s.r. spectrum at - 39.4 "C comprising a simple 1 : 1 : 1 triplet yielding a nitrogen hyperfine coupling constant of 2.660 & 0.01 3 mT.From our previous results,2 the latter value suggests that an arylalkoxy radical has been produced, since ultraviolet irradiation of 2,3,5,6-tetrachloronitrobenzene in tetrahydrofuran gives a radical having a three-line spectrum from a nitrogen splitting of 2.50 mT. The other solvents did not give readily assignable spectra. The improved equipment also held out the promise of observing monohydronitro radicals from other nitro compounds; however, no recognisable e.s.r. spectra could be obtained from nitrobenzene, para-nitrotoluene, 2,4,6-tri-t-butylnitrobenzene or nitromethane. These results caused us to limit the measurements made to 2,3,5,6- tetrachloronitrobenzene in propan-2-01. Table 1 lists the hyperfine coupling constants obtained by continuous irradiation at various temperatures.Reliable temperature coefficients cannot be calculated from these data but as agH definitely decreases with temperature its absolute sign is likely to be positive: we predicted5 that a temperature dependence should be detectable. Using the 1 kW ultraviolet lamp, a stationary-state radical concentration of 2.9 x dm3 mol-1 was obtained at -50 "C, more than one order of magnitude greater than observed in our previous inve~tigation.~ The laser gave a concentrationL. H. SUTCLIFFE 1469 Table 2. Second-order rate constants (k) for the decay of ArN(0)OH in propan-2-01 light source k/dm3 mol-1 s-l T/K 1 kW lamp 7.0 x 104 6.3 11.1 15.4 29.7 38.7 39.7 71.3 laser 5.2 x 105 6.4 10.4 18.0 22.6 214.5 222.9 233.1 241.3 249.8 260.3 280.3 291.3 236.9 259.1 268.7 293.1 313.5 Table 3.Energetics of the second-order decay of ArN(0)OH in propan-2-01 1 kW lamp k,,, = 1.34 x los dm3 mol-l s-' AG&, = 38 f 4 kJ mol-1 AH:,, = 17.5 & 2.0 kJ mol-l AS:,, = - 69 f 8 J mo1-1 K-l laser k,,, = 2.06 x los dm3 mol-l s-' AGZ,, = 37 t 5 kJ mol-l AH& = 14.2 f 2.5 kJ mol-l AS:,, = -77+9 J mol-l K-l of 7.4 x dm3 mol-1 at 20 "C. The decay of the radical was found to be accurately second order. Table 2 lists the second-order rate constants measured for the radical decay using the two photolytic techniques. Both sets of data gave good Arrhenius plots and the energetics derived from these are given in table 3. In our previous paper5 we placed an upper limit of lo6 dm3mol-l s-l (20 "C) on the second-order decay constant: it may be seen from table 3 that this was a good estimate. In view of the steric hindrance involved in the recombination of the radicals, it is not surprising that the rate constant is much smaller than that reported by Asmus et aL4 for the second-order decay of the monohydronitrobenzene radical.The energetic data obtained from the two kinetic methods are in good agreement. The appreciable negative entropy change is in accord with a radical-radical recom- bination mechanism. The laser method offered the possibility of measuring the rate of production of the radicals (< 0.02 s) but the time constant of the spectrometer was too long to allow the measurements to be made. The author is indebted to Dr K. U. Ingold, F.R.S. of the National Research Council, Ottawa for providing all the facilities used in this work.1470 E.S.R. STUDY OF THE MONOHYDRONITRO RADICAL D. J. Cowley and L. H. Sutcliffe, J. Chem. SOC., Chem. Commun., 1968, 201. D. J. Cowley and L. H. Sutcliffe, J. Chem. SOC. B, 1970, 569. M. McMillan and R. 0. C. Norman, J. Chem. SOC. B, 1968, 590. K. D. Asmus, A. Wigger and A. Henglein, Ber. Bunsenges. Phys. Chem., 1966, 70, 862. D. J. Cowley and L. H. Sutcliffe, Trans. Faraday SOC., 1969, 65, 2286. (PAPER 4/ 1646)

ISSN:0300-9599    

DOI:10.1039/F19858101467

出版商:RSC    

年代:1985

数据来源: RSC

摘要:

J. Chem. SOC., Faraday Trans. 1, 1985,81, 1471-1482 Kinetic Study of the Gas-phase Decomposition of the Tri fluor oace t yl Radical Effects of Temperature and Pressure upon the Rate Constants BY J. ALISTAIR KERR* AND J. PAUL WRIGHT Department of Chemistry, University of Birmingham, P.O. Box 363, Birmingham B15 2TT Received 1st October, 1984 The kinetics and pressure dependence of the decomposition of the trifluoroacetyl radical : CF,CO+M +CF,+CO+M ( 5 ' ) have been studied by generating the radicals from the selective photolysis of azomethane in the presence of l , l , 1-trifluoroacetaldehyde over the temperature range 338-41 7 K. The rate constants of the decomposition reaction (k,,) have been measured relative to the radical combination reactions : CH, + CF,CO 4 CF,COCH, (6) 2CH, + C,H,.(2) The rate constants (k5,), were found to be pressure-dependent over the pressure range 50-740 Torr, with SF, as a bath gas. The A factor of the radical decomposition reaction has been calculated from transition-state theory to be A? M 2.2 x lo1, s-l. Treatment of the pressure dependence of the rate constants (k5,), according to the RRKM theory of unimolecular reactions, yielded a hgh-pressure limiting activation energy E? = 83.1 kJ mol-l. Thus the present data are consistent with a high-pressure limiting Arrhenius expression : log (kP/s-') = (1 3.34 f 0.8) - [( 10 000 & 1000)/2.303 TI where the error limits are estimates of the accuracies of the Arrhenius parameters. with kinetic data on the decomposition reactions of the CH,CO and C,H5C0 radicals. The results are discussed in relation to existing data on the CF,CO radical and are compared Kinetic and thermochemical data on acyl radicals are important in their own right and have become increasingly in demand as the need for information on radical reactions grows with the development of kinetic modelling of complex systems such as combustions and atmospheric chemistry.Quantitative information on the decomposition reactions of acyl radicals is fragmentary and has a history of contro~ersy.l-~ While the rate constants for the decompositions of the a ~ e t y l ~ - ~ and propionyl8-l0 radicals appear to be reasonably well established and in line with current theories of chemical kinetics, for other acyl radicals few rate constants have been confirmed.The trifluoroacetyl radical is a case in point, where the sole quantitative study'' reports an A factor approximately three powers of ten lower than the minimum A factor required by transition-state theory. In addition the observed rate constantsll were found to be independent of pressure up to a total pressure of ca. 600 Torr,? in t 1 Torr = 101 325/760 Pa. 14711472 DECOMPOSITION OF THE TRIFLUOROACETYL RADICAL contrast to the predictions of unimolecular reaction rate theory and the observed pressure dependences of the CH,CO and C,H,CO radical decomposition reactions. Here we report a further investigation of the kinetics of the decomposition of the CF,CO radical by a method which has yielded consistent data for the CH,C012 and C,H,C08 radicals. EXPERIMENTAL APPARATUS AND PROCEDURE Runs were carried out in a 159 cm3 cylindrical quartz reaction vessel contained in an electrically heated furnace and controlled to better than +1 K.Reactant pressures were measured with a mercury manometer attached directly to the cell. Radiation of wavelength centred around 366 nm was isolated from a 250 W medium-pressure mercury arc by means of Corning filters no. 7-37 and 0-52. Under these conditions there was no photolysis of the trifluoroacetaldehyde. ANALYSIS CH,, CO and N, were pumped off at liquid-air temperature with a Topler pump and analysed by gas chromatography. The column consisted of 2.25 m of 30-60 mesh 5 8, molecular sieve and the carrier gas was H, with a thermal-conductivity detector. Temperature programming was carried out with the column initially at 323 K for 4 min followed by an increase of 30 K min-l to 383 K. Under these conditions CO had a retention time of 6 min.C,H, was pumped off at 123 K from a Ward still via the Topler pump. Gas-chromatographic analysis was carried out on a 1.75 m column packed with 100-120 mesh Porapak T and operated isothermally at 353 K with' flame ionization detection. The remainder of the products and reactants were then condensed into the injection loop of the gas chromatograph and analysed on the same column as for C,H, but this time with temperature programming and flame ionization detection. The column was operated at room temperature for 4 min followed by a temperature rise of 35 K min-' to 458 K. The retention time of 1, 1 , 1-trifluoroacetone was 23 min.Quantitative calibrations for the products CO, C,H, and CF,COCH, were carried out with authentic samples of these materials. MATERIALS Azomethane was prepared from sym-dimethylhydrazine dihydrochloride (Aldrich Chemical Co.) by the method described by Renaud and Leitch.', 1,1,1 -Trifluoroacetaldehyde was obtained as the hydrate (Lancaster Syntheses Ltd) and dehydrated with a stirred mixture of P,O, and concentrated H,SO, maintained at 363 K. SF, (Cambrian Gases, C.P. grade) was used without further purification. RESULTS The proposed mechanism, to account for the formation of the products N,, CH,, CO, C,H6 and CF,COCH, formed by the selective photolysis of CH,N=NCH, in the presence of CF,CHO, is as follows: CH,N=NCH, + Av + X H , + N, (1) 2CH, -+ C2H6 ( 2 ) (3) CH, + CF,CHO -+ CH, + CF,CO CF,CO + M a CF,CO* + M CF,CO* + CF, + CO (4, -4) ( 5 )J.A. KERR AND J. P. WRIGHT 1473 CH, + CF3C0 -+ CF,COCH, (6) (7) + (CH,),NNCH, (8) (9) (10) CH, + CH,N=NCH, -+ CH4 + CH,N=NCH, CH,N=NCH, + R -+ RCH,N=NCH, (CH,),NNCH, + R -+ (CH,),NN(CH,)R. A steady-state treatment of the activated CF,CO* radicals gives d[CH,CO*]/dt = k, [CF,CO] [MI - k-, [CF,CO*] [MI - k5 [CF,CO*] = 0. (I) We can also write the rate equations: RCF3COCH3 = k6 [CH3] [CF,COI Rco = k,[CF,CO*] = k2 [CH3I2 where Rx is the rate of formation of product X. Substitution of these three rate equations into eqn (I) and rearranging yields RCF3COCH3/(RCOkC2H,,) = J? = (k-4 k 6 ) / ( k 5 k 4 k . t , ) + k 6 / ( k 4 k!? (11) Thus a plot of the rate function, F, against 1 /[MI, a classical Hinshelwood-Lindemann plot, should be linear.In this treatment [MI is taken to be the initial total concentration of all reactants, on the basis that the reactions proceed to low percentage conversions. From eqn (11) it follows that as [MI +a, i.e. I/[M] -+ 0 F = (kP4 k,)/(k5 k, k i ) = k , / ( k p #) and that k,/(kp k!) can be calculated from the intercept of the plot of Fagainst 1 /[MI, where k? is the high-pressure limit of the decomposition reaction CF,CO+M +CF,+CO+M. ( 5 ' ) Absolute values of k p were obtained by taking k, = 2.2 x 1O1O dm3 mol-l s-l independent of temperat~re'~ and by making the reasonable assumption that k , = k,. The experimental results for the temperatures 338, 377, 395 and 417 K are shown in table 1.The first point to note is that F displays distinct pressure dependence, for instance at 395 K, Pdecreases by a factor of 3.3 over the pressure range 26-658 Torr. This is good evidence for the pressure dependence of the CF,CO decomposition reaction. It was not possible to carry out experiments below 338 K owing to the ready polymerization of the a1deh~de.l~ At the same time the temperature range could not be extended beyond 417 K as the yields of CF,COCH, became too small for accurate measurement. The apparatus was not designed to operate at total pressures exceeding 1 atm. The range of experimental conditions, which we recognize to be limited, was the maximum obtainable with the present system. The Hinshelwood-Lindemann plots of the experimental data are reasonably good straight lines, which by least-mean-squares analyses yield the following values of k? : 338 377 395 417 4.05 42.8 117 376.T / K k? /s-l These in turn correspond to the following Arrhenius equation: k? = 1 .O x loll exp (- 8080/ T ) s-lTable 1. Rate data from the photolysis of azomethane in the presence of trifluoroacetaldehyde c, P P F = RCF3COCH3 4 RCO @C2Hs P(CHdQ2 PCF3CHO Ptotal R C O RCzHs R C F 3 C O C H 3 /Torr /Torr /Torrs mol dm-3 s-' /lop8 mol dm-3 s-' mol dm-3 s-' /lo3 mol-: dm2 st 10 17 14 23 17 20 20 30 29 44 27 42 22 39 27 27 24 32 21 35 30 41 34 35 10 16 12 22 10 20 18 28 13 25 18 36 27 40 30 44 31 47 30 41 28 42 27 41 26 35 32 40 32 42 34 55 31 56 29 44 27 37 37 50 73 107 197 287 324 380 527 624 26 34 42 46 48 54 67 74 78 105 154 208 254 318 378 497 613 739 0.28 0.54 0.30 0.83 1.30 0.87 0.65 1.15 1.10 1.42 1.29 1.78 0.47 0.70 0.71 1.03 0.49 1.24 1.82 1.96 1.82 1.55 1.36 1.58 1.19 1.91 1.26 0.63 1.02 0.94 T = 338 K 0.26 0.56 2.60 0.3 1 1.35 0.77 0.84 1.09 0.84 0.28 0.87 0.80 0.82 2.36 0.63 3.0 1 1.42 2.38 2.08 1.98 6.23 1.18 8.18 2.95 2.04 0.68 7.63 7.95 6.73 3.52 T = 377 K 0.75 2.06 2.39 2.22 6.85 3.43 2.46 4.57 3.57 2.62 4.44 5.37 0.35 0.80 0.4 1 I .09 0.41 1.20 1.60 1.61 2.68 0.89 1.71 1.09 0.7 1 0.63 1.29 0.64 9.3 1 6.28 52.5 51.0 49.4 48.0 45.4 44.9 41.3 38.1 35.4 34.9 36.9 33.7 8.22 7.44 7.28 6.10 7.02 6.27 6.10 5.83 5.90 5.29 4.40 4.02 4.18 3.99 3.71 3.60 3.52 3.56 cl 76 9 c, 9 s r10 11 13 14 18 20 27 27 37 30 29 25 39 17 25 36 17 9 10 12 18 15 24 39 20 26 40 36 38 46 44 16 19 20 23 32 30 44 58 40 35 40 33 55 22 48 68 31 18 20 30 24 35 45 40 41 40 61 48 47 48 52 26 30 33 37 50 50 71 85 112 171 171 217 354 389 497 528 658 27 30 42 42 50 69 79 128 181 279 316 434 478 624 0.89 0.72 0.69 0.97 1.44 1.76 2.89 3.08 2.28 1.18 0.71 1.26 0.82 1.10 0.90 1.12 0.74 0.68 0.79 1.07 1.02 1.15 1.81 1.80 1.29 1.15 1.32 1.21 0.98 2.22 1.90 T = 395 K 0.33 0.32 0.5 1 0.70 1.04 0.62 0.55 0.43 0.98 0.65 1.07 2.89 2.38 2.47 4.53 2.55 3.86 1.79 3.34 1.30 0.65 3.26 3.40 0.75 1.45 5.28 3.23 5.22 3.63 5.80 2.49 T = 417 K 0.20 0.14 0.17 0.25 0.38 0.40 0.54 0.46 0.45 0.17 0.1 1 0.28 0.16 0.27 0.24 0.27 0.17 0.23 0.34 0.23 0.14 0.3 1 0.43 0.20 0.14 0.17 0.14 0.20 0.10 0.23 0.11 3.91 3.43 3.45 3.08 2.58 2.89 2.52 2.23 1.99 1.79 1 S O 1.31 1.26 1.56 1.25 1.51 1.17 2.53 2.35 1.88 1.70 1.49 1.29 1.28 0.901 0.643 0.590 0.724 0.536 0.430 0.367 4 co1476 DECOMPOSITION OF THE TRIFLUOROACETYL RADICAL At the mean temperature of these experiments the universal frequency factor for a unimolecular reaction (ekT/h) corresponds to 1.9 x 1 013 s-l and consequently the derived value of A? = 1.0 x loll s-l imples a large negative entropy of activation.Similarly low A factors were originally reported for the CH,CO decompositions but these have been shown to be wrong. The problem in studies of the present type is that the experimental rate constants have been measured too far from the high-pressure limiting conditions and accurate extrapolations are difficult to make. Two alternative extrapolation techniques were applied.16 In the first a plot of I l k against 1 / f i was made and again the intercept yields a value of kCC, which often corresponds to an overestimate.16 The present data for temperatures of 338, 377 and 395 K were treated in this way but the data for 417 K gave a negative intercept and were therefore neglected.The derived values of k" are as follows: 338 377 395 4.8 68.6 273 T / K k p /s-l corresponding to the Arrhenius equation k? = 4.3 x 10l2 exp(-9320/T) s-l. The second extrapolation procedure involves plotting 1 / k against 1 / p a , where a is an adjustable parameter with a value between 0 and 1. a is optimised to give the best straight-line fit of the data. From the data of table 1, k? values derived in this way are as follows: 338 377 395 417 T / K a k? 1s-l 4.8 58.6 155 582.0.5 0.6 0.7 0.8 These yield the following Arrhenius equation, when treated by a least-man-squares analysis : k? = 4.0 x loll exp ( - 8520/T) s-l. It is clear that while these alternative extrapolations offer some improvement on the Hinshelwood-Lindemann procedure, none of the treatments yields an acceptable value of the A factor for the radical decomposition reaction. Accordingly a different approach was applied. The high-pressure limiting A factor for the radical decomposition reaction was calculated from transition-state theory, and the pressure dependence of the rate constants was treated by the RRKM theory of unimolecular reactions to derive the high-pressure limiting activation energy. For the CH,CO and C,H,CO radical decomposition reactions the experimental A-factors5-l0 are close to the value of (ekT/h) indicating that the structure of the activated complex is close to that of the reactant radical.These structures can be represented as R-C +, R..... c+R + co. 1. 0 \\ 0 reactant activated complex The main structural changes contributing to A S in this case involve changes in the vibrational modes. The vibrational assignments for the CF,CO radical have been made from the experimental frequencies1' in the aldehyde CF,CHO. The activated complex was assumed to involve a loosening of the C-C=O bending frequency andJ. A. KERR AND J. P. WRIGHT 1477 -1.0 0 I .o 2.0 3.0 log,, (plTorr) Fig. 1. (-) Calculated RRKM fall-off curve (Forst program) optimized to experimental rate constants at 377 K (0). a slight tightening of the C=O stretching frequency, in accord with the structures proposed above.The assigned vibrational frequencies are listed in the Appendix. The A factor for reaction ( 5 ’ ) was then calculated from the transition-state theory equation log (A,./s-l) = log (ek/h) + log (T/K) + AS/2.303R in which the entropy of activation, A S = 0.29 J mol-1 K-l, was calculated from the changes in vibrational frequencies and the tabulation of assigned entropie~,~ and the other terms have their usual meaning. For the mean temperature of these experiments, 337 K, this corresponds to A,. = 2.2 x 1013 s-l. The development and application of the RRKM theory of unimolecular reactions have been described in detail by Robinson and H~lbrook.~ The calculations are most readily achieved with the aid of computer programs. For the present purposes two such programs were applied, a Forst program and a Marcus program.The Forst program requires a knowledge of the Arrhenius parameters for the reaction as well as the usual molecular properties. The Marcus program on the other hand, calculates the A factor from the vibrational assignments of the reactant and activation complex. The input data required for the computer program are listed in the Appendix. Taking the previously calculated value A,. = 2.2 x 1013 s-l, the activation energy was optimised to the fall-off data at the experimental mid-point temperature of 377 K by the use of the Forst program. This resulted in a value of E5! = 83.1 kJ mol-l. The derived pressure fall-off curve is compared with the experimental data in fig.1. The Marcus program was then used to treat the data at all four temperatures and the corresponding fall-off curves are shown in fig. 2 and 3. The best fits were obtained at 377 and 395 K. At 338 and 417 K the calculated curves were an underestimate and overestimate, respectively, in relation to the experimental data, which could be interpreted as indicating a systematic error. Better fits of the data could be derived by adjustments to the estimated A factor, but this hardly seems justified. We conclude that the present data for reaction (5’) are consistent with the Arrhenius equation log (k? /s-’) = (1 3.34 f 0.8) - [( 10000 1000)/2.303 r]1478 DECOMPOSITION OF THE TRIFLUOROACETYL RADICAL I .5 1.0 .4 I VY 1 0.5 2 M - O o - I I I I I I - - I I I I .o 0 I .o 2.0 3.0 h 2.0 1 v) Y, 2 I .o - M 0 log,, (plTorr) Fig, 2.(-) Calculated RRKM fall-off curves (Marcus program) compared with experimental rate constants at 377 (0) and 395 (m) K. 0 4 M c1 log,, (plTorr) Fig. 3. (-) Calculated RRKM fall-off curves (Marcus program) compared with experimental rate constants at 338 (0) and 417 (w) K. in which the error limits are estimates of the overall accuracies of the derived A factor and temperature coefficient and reflect the problems inherent in the interpretation of data obtained in the pressure fall-off region. DISCUSSION KINETIC STUDIES OF THE CF,CO DECOMPOSITION REACTION Amphlett and Whittle" obtained kinetic data on the CF,CO decomposition reaction CF,CO --* CF, + CO ( 5 )J. A. KERR AND J.P. WRIGHT 1479 relative to the reactions CF,CO + X, -+ CF,COX + X (1 1) where X, = C1, or Br,, by photolysing the halogen in the presence of CF,CHO. Over the temperature range 300-443 K they determined log ( k , / ~ - ' ) = 10.42 - (5030/2.303T) and in view of their assumed values of k,, it was recommended that A , and E5 be regarded as upper and lower limits, respectively. This paper also incorporates some earlier data of Tucker and Whittle1* on reaction (9, derived from the photolysis of (CF,),CO in the presence of Cl,. A comparison of the rate constants at 377 K reveals a value of k, = 4.2 x lo4 s-l as measured by Amphlett and Whittle,' and a value of k p = 66.5 s-l from the present data. At present we have no explanations of this very large discrepancy between the two studies.One could argue that the values of the rate constants (kll) of the reference reactions assumed by Amphlett and Whittle might be too high, but it seems unlikely that any discrepancy in the k,, values would amount to as much as a factor of 600. The abnormally low A factor reported by Amphlett and Whittle was at the time of publication compatible with the existing literature value of the A factor for the CH,CO radical decomposition, which was subsequently shown to be much closer to (ekT/h). Amphlett and Whittle considered the possibility that their rate constants (k,) had been measured in the pressure fall-off region. The bulk of their experiments, performed over the pressure range 20-80 Torr, showed no variation with pressure, as did a reaction carried out in the presence of 550 Torr of added N,.In view of the established pressure dependences of the CH,CO and C,H,CO decompositions, together with the present finding of a distinct pressure dependence of the CF,CO decomposition reaction, the lack of such an effect in the experiments of Amphlett and Whittle is puzzling. McIntosh and Porterlg attempted to determine the kinetic stability of the CF,CO radical from flash photolysis of (CF,),CO alone and in the presence of Br,. Their derived data for reaction (5) were inconclusive, however, owing to the likely generation of excited CF,CO radicals from the primary photochemical dissociation. One other set of relevant experiments involving CF,CHO have been reported by Loucks et aZ.,,O who pyrolysed di-t-butyl peroxide in the presence of the aldehyde over the temperature range 390-441 K.CF,COCH, was observed as a product and was ascribed to the following reaction sequence : (12) CH, + CF,CHO -+ CH,(CF,)CHO CH,(CF,)CHO -, CF,COCH,+H. (13) The occurrence of the reaction CH, + CF,CO -+ CF,COCH, (6) in this system was discounted on the grounds of the lack of thermal stability of the CF,CO radical at these temperatures. Small amounts of the alcohol CH,(CF,)CHOH were detected, which is good evidence for the occurrence of reaction (12) followed CH,(CF,)CHO + CF,CHO -+ CH,(CF,)CHOH + CF,CO. by1480 DECOMPOSITION OF THE TRIFLUOROACETYL RADICAL The exclusive decomposition of the alkoxy radical according to reaction (13) in preference to the alternative processes CH,(CF,)CHO + CH,+CF,CHO -+ CF, + CH,CHO is questionable on thermochemical grounds.In addition, recent studies2l9 22 of the decompositions of the analogous alkoxy radicals, CH, (CF,),Co and CF, (CH,),Co, have shown that decomposition in each case occurs exclusively via CF, loss. A further argument against reactions (1 2) and (1 3) forming the exclusive route to CF,COCH, formation in the Loucks et aL20 system comes from their observation that the yields of this product were 4-5 times larger than the yields of H,. The latter would be expected to be the major product of the H atoms in the presence of the labile CH of the parent aldehyde. Even if reaction (12) did occur to some extent, the lack of equality between CF,COCH, and H, implies an additional source of CF,COCH, in this system, which we contend is reaction (6).The Arrhenius parameters obtained by Loucks et CZZ.~O for the proposed reaction (1 2) k,, = 5.0 x lo6 exp (- 2010/T) dm3 mol-' s-I are highly suspect. Both the A factor and the activation energy would appear to be seriously underestimated on the basis of transition-state theory and in relation to recent related data for alkyl-radical additions to carbonyl groups.22 Unfortunately Loucks et aI.,O provide no analytical data on the rates of formation of CO in their study, which would have enabled a test of their system in terms of a mechanism involving reactions ( 5 ) and (6). It is possible, however, to treat our data according to a mechanism containing reactions (1 2) and (1 3). This kinetic analysis is similar to those which we have recently used for CH, additions to CF,COCF,21 and to CF,COCH,.22 If reactions (12) and (13) are the only sources of CF,COCH, formation then for a fixed temperature a plot of RCF,COCH,/R~C2H,, against [CF,COCH,] should be linear and pass through the origin.When treated in this fashion, the present data do not show such linear dependence. While this does not rule out the formation of some of the CF,COCH, by reactions (12) and (13), it disproves the carbonyl- addition mechanism as the exclusive route to this product. COMPARISON OF CF,CO, CH,CO AND C,H,CO DECOMPOSITIONS It is interesting to compare the present results for the decomposition of the CF,CO radical with data for the related CH,CO and C,H,CO radical reactions. As seen from table 2 the rate constants at 377 K reveal that the sequence of kinetic stabilities runs CF,CO > CH,CO > C,H,CO, which is a reflection of the activation energies for the decomposition reactions since the A factors are effectively equal.On the other hand, in its pressure fall-off behaviour the CF,CO radical is much closer to C,H,CO than to CH,CO, as shown by the 4 values, which are the pressures at which k , falls to one-half of its high-pressure value. A, vaiues are determined by the critical energy of the reaction together with the number of effective oscillators of the molecule. Thus the large difference between the A, values for CH,CO and CF,CO is accounted for by the increased number of effective oscillators arising from the F atoms in CF,CO, together with the larger Em for CF,CO.In the case of C2H5C0 and CF,CO the greater number of effective oscillators for C,H5C0 compared with CF,CO is offset by the smaller activation energy for the decomposition of C,H,CO compared with CF,CO.J. A. KERR AND J. P. WRIGHT 148 1 'Table 2. Comparison of the kinetics of the decompositions of the radicals CF3C0, CH3C0 and C,H,CO radical Em/kJ mol-l log ( A w l s - ' ) km/s-l a p:/Torra CF3C0 83.1 13.3 6.0 x lo1 56 CH,CO 72.0 13.5 3.3 x 103 1 135b C,H,CO 61.5 13.3 6.0 x 104 216c a At 377 K. Extrapolated from the data of Watkins and Word7 and of O'Neal and Extrapolated from the data of Watkins and Thompsonlo and of Kerr and Lloyd.B B e n s ~ n . ~ ~ CONCLUSIONS The present kinetic study of the CF,CO decomposition reaction suggests a rate constant given by k? z 2.2 x lo1, exp (- 10000/T) s-l.This result is consistent with literature data for the kinetics of the CH,CO and C,H,CO decomposition reactions but is seriously at variance with previously published data," which indicate rate constants ca. 600 times greater than those corresponding to the above Arrhenius equation. In view of this large unexplained discrepancy, it would be desirable to obtain further data on reaction ( 5 ' ) by an independent study. We are presently attempting to investigate the kinetics of the system CF,+CO+M -", CF,CO+M by generating CF, from the photolysis of CF,N=-NCF, in the presence of CO. We thank Drs K. A. Holbrook and R. Walsh for helpful discussions on the RRKM calculations and for the use of their computer programs.J. A. Kerr and A. C. Lloyd, Q. Reu., 1968, 22, 549. S. W. Benson and H. E. ONeal, Kinetic Data on Gas Phase Unimolecular Reactions, NSRDS-NBS 21 (U.S. Government Printing Office, Washington D.C., 1970). P. J. Robinson and K. A. Holbrook, Unimofecufar Reactions (Wiley-Interscience, London, 1972). K. A. Holbrook, Chem. SOC. Rev., 1983, 12, 163. H. M. Frey and I. C. Vinall, Int. J. Chem. Kinet., 1973, 5, 523. L. Szirovicza and R. Walsh, J. Chem. SOC., Faraday Trans. 1, 1974, 70, 33. K. W. Watkins and W. W. Word, Int. J. Chem. Kinet., 1974, 6, 855. P. Cadman, C. Dodwell, A. F. Trotman-Dickenson and A. J. White, J. Chem. SOC. A , 1970, 2371. * J. A. Kerr and A. C. Lloyd, Trans. Faraday SOC., 1967, 63, 2480. lo K. W. Watkins and W. W. Thompson, Int. J. Chem. Kinet., 1973, 5, 791. l1 J. C. Amphlett and E. Whittle, Trans. Faraday SOC., 1967, 63, 80. l2 J. A. Kerr and J. G. Calvert, J. Am. Chem. SOC., 1961, 83, 3391. l3 R. Renaud and L. C. Leitch, Can. J. Chem., 1954,32, 545. l4 A. Shepp, J. Chem. Phys., 1956, 24, 939. N. L. Arthur and T. N. Bell, Aust. J. Chem., 1965, 18, 1561. l6 I. Oref and B. S . Rabinovitch, J. Phys. Chem., 1968, 72, 4488. R. E. Dodd, H. L. Roberts and L. A. Woodward, J. Chem. SOC., 1957,2783. l8 B. G. Tucker and E. Whittle, Trans. Faraday SOC., 1965, 61, 866. J. S. E. McIntosh and G. B. Porter, Trans. Faraday SOC., 1968, 64, 119. L. F. Loucks, M. T. H. Liu and D. G. Hooper, Can. J. Chem., 1979, 57, 2201. 49 FAR 11482 DECOMPOSITION OF THE TRIFLUOROACETYL RADICAL 21 R. M. Drew and J. A. Kerr, Int. J. Chem. Kinet., 1983, 15, 281. 22 J. A. Kerr and J. P. Wright, Znt. J. Chem. Kinet., 1984, 16, 1327. 23 H. E. O’Neal and S. W. Benson, J. Chem. Phys., 1962, 36, 2196. (PAPER 4/ 1698) APPENDIX TABLE Vibrational assignmentsa CF,CHO CF,CO CF,COI C=O str. C-F str. C-F str. C-F str. C-C str. CF, def. CF, def. CF, def. C-C=O bend CF, rock CF, rock torsion 1770 1305 1200 1140 957 840 710 530 430 320 265 184 2100 2100 1305 1305 1200 1200 1140 1140 370 rcb 840 840 710 710 530 530 220 100 215 215 180 180 150 150 a Collision diameter of SF,, CT = 5 x cm; collision efficiency of SF, = 0.5. Reaction coordinate.

ISSN:0300-9599    

DOI:10.1039/F19858101471

出版商:RSC    

年代:1985

数据来源: RSC

摘要:

J. Chem. Soc., Faraday Trans. 1, 1985, 81, 1483-1494 Deuterium Isotope Effect in Concentrated Aqueous Solutions A Potentiometric and 13C Nuclear Magnetic Resonance Study of Acid Dissociation Constants BY STEPHEN P. DAGNALL, DAVID N. HAGUE,* MICHAEL E. MCADAM AND ANTHONY D. MORETON University Chemical Laboratory, Canterbury, Kent CT2 7NH Received 8th October, 1984 The acid dissociation constants have been determined by potentiometric titration in H,O and D,O for the zinc(I1) complexes of 2,2’-ethylenedi-iminobis(ethylamine) (trien), 2,2’- trimethylenedi-iminobis(ethy1amine) (2,3,2-tet) and 2,2’,2”-nitrilotris(ethylamine) (tren) and by ‘“C n.m.r. spectroscopy in D,O for ethylenediamine (en;2), glycine (gly;2), L-alanine (ala), acetic acid (ace) and N-(2-aminoethyl)propane- 1,3-diamine (2,3-tri ;3).The equilibrium deuterium isotope effects are similar, pKa(D,O) - pKa(H,O) = 0.63 & 0.07. This result supports the treatment of the protonated forms of polyamines as ‘hybrid’ ions in respect of their contribution to ionic strength and indicates that the coordination number of Zn” in the low-pH forms of the three complexes is at least five. It is generally agreed that a major reason for the importance of zinc in enzyme chemistry is its ability to adopt a coordination number of 4, 5 or 6 and to change its coordination geometry relatively easily. The range of techniques that can be used to determine the coordination number of complexed Zn“ is severely limited by its d10 configuration, and the only ones which have been used at all widely are X-ray single-crystal and powder diffraction.This unavoidable dependence on solid-state data has seriously impaired mechanistic discussions on the role of zinc in enzymatic reactions, and much use has been made of evidence from analogous reactions with Co” and other species with suitable spectral properties. We have become interested in the possibility of using the 13C n.m.r. spectra of coordinated organic ligands to explore the geometry of Zn” complexes in aqueous solution and will report on this in future papers. Most of these complexes are labile on the n.m.r. timescale and the observed chemical shifts are weighted averages of contributions from the various species present, namely one or more complexes and the uncomplexed form(s) of the ligand(s). A convenient way of separating these contributions is to study the pH dependence of the n.m.r.spectra but a detailed analysis requires reliable acid dissociation constants for the ligands and, in many cases, the coordinated metal ion itself. Since for technical reasons it is necessary to use concentrations in the decimolar range and D,O as the solvent, the thermodynamic data in the literature must be adjusted for the deuterium isotope effect (DIE) and possibly for ionic strength. The present paper is concerned with these two adjustments and reports data on five representative free ligands and three zinc complexes. At low ionic strength I, the activity coefficient yi of the ith ion carrying a charge zi is given by the Debye-Huckel limiting law, log yi = -Az,ZIi, where the constant A depends on the temperature and solvent permittivity.Several empirical modifications 49-2 14831484 DIE STUDY OF pKa have been derived for aqueous solutions of ionic strength 2 0.01 mol dmP2, of which the Daviesl equation is well established. There is good agreement between measured and calculated (mean ionic) activity coefficients for electrolytes involving simple ions, but yi for large and complex ions are often more difficult to predict.2 Many of the ligands we use are bi- or multi-dentate, and their metal-binding sites are either charged (e.g. carboxylate groups) or become charged on protonation within the pH ranges of interest (e.g. amino groups). The free ligands may therefore carry a large, and varying, net charge and it is important to decide on the contribution they make to the ionic strength and consequently to the apparent acid dissociation constant.There have been many investigations into the effect of changing the solvent from H 2 0 to D,O on the dissociation constants of weak acids [see, for example, ref. (3)]. Early results indicated4* that it was larger the weaker the acid and the existence of a precise linear relationship between ApK[ = pKa(D20) - pKa(H,O)] and pKa(H20) has been demonstrated for several series of related acids [see, for example, ref. (6)-(9)]. In other cases of similar acids there is no such increase of ApK with pKa(H20) [see, for example, ref. (7), (10) and (1 l)]. We report the pKa values determined in D 2 0 at moderately high ionic strengths ( I = 0.1-0.6 mol dm-3) of the ligands ethylenediamine (en), glycine (gly), L-alanine (ala), acetic acid (ace) and N-(2-aminoethyl)propane- 1,3-diarnine (2,3-tri), and of the zinc complexes of 2,2'-e t hylenedi-iminobi s( e t h ylamine) (t rien), 2,2'- trime t h ylene- di-iminobis(ethy1amine) (2,3,2-tet) and 2,2',2"-nitrilotris(ethylamine) (tren).In the case of the ligands, the data were obtained from the titration of their 13C n.m.r. chemical shifts and the observed pKa are compared with values (in H20) from the literature. For the zinc complexes the data were obtained by potentiometric titration and, since no literature values are available for comparison, the experiments were repeated in H20. (Differences in the 13C n.m.r. shifts for the protonated and deprotonated forms of the complexes are rather sma11.12) The DIE is found to be similar for the 12 individual pKa studied (0.63+0.07), a result which supports the treatment of the protonated forms of polyamines as 'hybrid ' ions13 and indicates that the coordination number of Zn" in the low-pH forms of these three complexes is at least 5.log yi = - Az;[z:/( 1 + Zi) - 0.311 (1) EXPERIMENTAL The amines were purified by distillation as follows and their purity monitored by 13C n.m.r. in conjunction with analytical g.1.c.: en (B.D.H.; 116-1 17 "C at 760 Torr), 2,3-tri (Aldrich; 60-64 "C at 1 Torr), trien (from the sulphate, Baker; 120-130 "C at 0.005-0.1 Torr), 2,3,2-tet (Eastman; 100-102 "C at 0.05 Torr), tren (isolated from technical-grade trien as the hydr~chloride;'~ 127-128 "C at 0.01 5 Torr).Acetic acid, 1,4-dioxane, glycine (Fisons A.R.) and L-alanine (Koch-Light puriss.) were used without further purification. Solutions of the zinc complexes were made up from aqueous Zn(NO,),, which was prepared from ZnO (Fisons A.R.) and a slight excess of concentrated nitric acid and estimated15 against EDTA. Solutions (generally 0.2moldm-3) were made up in D,O (B.O.C. Prochem; 99.8 atom%) or triply distilled water; those for the n.m.r. experiments contained 0.2 mol dmP3 1,4-dioxane. The pH/pD measurements were made on a Radiometer PHM 26C with GK 2321C semimicro dual electrode, and Na+-ion corrections were applied at high pH according to the manufacturer's nomogram (in the case of the measurements in D,O, to the meter reading) (see Results). The electrodes were standardised against N.B.S.phosphate and borax buffers,16 which were used within 3 weeks of their preparation. The pD was calculated using the empirical relationship13 pD = (meter reading) + 0.40; in the n.m.r. experiments it was adjusted with concentrated HNO, and/or a saturated solution of NaOH in D,O.S . P. DAGNALL, D. N. HAGUE, M. E. MCADAM AND A. D. MORETON 1485 The potentiometric titrations were performed under a C0,-free atmosphere in a thermostatted (20.0 f 0.1 "C) glass thimble fitted with a magnetic stirrer using standard techniques." Generally, NaOH/NaOD (1 .OO mol dm-3) was delivered from a 1 cm3 syringe into 5.00 cm3 of zinc/amine solution (0.20/0.22 mol dmp3), although confirmatory titrations were run using [zinc/amine] = 0.20/0.21 mol dmp3 and [NaOH] = 0.3 and 0.4 mol dm-3.The n.m.r. spectra were recorded at ca. 20 "C with a JEOL FT-100 spectrometer in the proton-noise-decoupled mode at 25.15 MHz. Typically, 4000 data points were accumulated over a frequency range of 2 kHz with 500 scans. The chemical shifts were measured relative to internal 1,4-dioxane but the data are quoted on the 6 scale. In several experiments designed to confirm the lack of influence of typical solutes and pD on the signal of 1,4-dioxane, external Me,Si was used as a second standard: the 6 value of dioxane was always found to be 67.71 and we have therefore used this value throughout. (The addition of this amount of 1,4-dioxane appearslS to have no significant effect on the pK, of aliphatic polyamines.) In the cases of gly, ala and ace, the collection of data was restricted to the CH, (n = 1-3) carbons owing to the comparatively low intensity (from long and lack of n.0.e.) and large chemical shift of the carboxylate resonances.RESULTS POTENTIOMETRIC TITRATIONS The measured pH or pD values for H,O and D20 solutions of Zn(trien)2+, Zn(2,3,2-tet),+ and Zn(tren),+ to which varying amounts of 1.00 mol dm-3 NaOH or NaOD, respectively, have been added are indicated by the solid circles in fig. 1 . The open circles are the same experimentally determined points to which the appropriate sodium-ion correction has been applied. The titration curves have been computed from the data by the following method, which is based on standard complexometric techniques. l9 At a particular point on the titration curve for an H 2 0 solution the measured pH will be simply related through K,, the ionic product of water, to the total amount of NaOH which has been added less the amount of OH- which has reacted.If we assurne2O that only 1 : 1 complexes are formed between Znit and these ligands L, then the only species with which OH- can react are the small amount of (protonated) excess ligand [added to prevent precipitation of Zn(OH), at moderately alkaline pH], free H,+, and the zinc complex. We shall discuss the formulation of these complexes in more detail in future papers (see also the Discussion) but for the present we shall refer to the low-pH form as ZnL(OH,),+ and the high-pH form as ZnL(OH)+. On this basis, the main reaction occurring during the titration is ZnL(OH2)2+ +OH- -+ ZnL(OH)+ + H20 whose equilibrium position can be characterised by the acid dissociation constant K, of the zinc complex ZnL(OH,),+.It would not alter the titration analysis if the OH- ion were to add to the low-pH form of the complex (ZnL2+ + OH- + ZnLOH+) rather than remove a proton from it, although it would, of course, modify the interpretation of the result. Incidentally, the concentration of the monoprotonated form of the trien complex2o is negligible over the pH range covered by our titrations. A preliminary value of K, was assumed, based on the inflexion point of the experimental titration curve. This was used to calculate, by means of a quadratic in the concentration of unprotonated free ligand, the concentrations of all L-containing species at the pH corresponding to one of the experimental points.From the calculated value of [ZnL(OH)+] and the actual volume of NaOH used, a new value of K, was determined and the cycle repeated until two consecutive values of K, differed by < 1 % . This iterative procedure was performed at the pH of each experimental point and a1486 DIE STUDY OF p& 1.00 - rn E \ W 8 0.75- 4 0 I I 1 I I I I 8 .O 8.5 9.0 9.5 10.0 10.5 11.0 11.5 pH or pD 9.0 9.5 10.0 10.5 11.0 11.5 12.0 12.5 13.0 pH or pD pH or pD Fig. 1. Potentiometric titrations of complexes in H,O and D,O. NaOH or NaOD (1.00 mol dm-3) was added to 5.00 cm3 of a solution containing 0.20 mol dm-3 zinc and 0.22 mol dmP3 amine (see text). (a) Zn(trien), (b) Zn(2,3,2-tet) and (c) Zn(tren).S.P. DAGNALL, D . N. HAGUE, M. E. MCADAM AND A. D. MORETON 1487 Table 1. Measureda pK, values for zinc complexes in H,O and D,O complex P G PKh Zn(trien) 9.07 9.74 Zn(2,3,2-tet) 10.30 10.91 Zn( tren) 10.59 11.21 a 20.0f0.1 “C, I = 0.4 mol dm-3; the estimated errors are fO.l (but see text). weighted average taken of the resulting Ka values. Allowance was made for dilution accompanying each addition of NaOH and, by means of eqn (l), for the the varying ionic strength and its effect on the single-ion activity coefficients of charged species. The same procedure was adopted for the D 2 0 titrations using pD values determined by adding 0.40 to the meter readings. The ‘best’ value of Ka was then used to generate the volume against pH(pD) curves shown in fig. 1. For the calculations, the following literaturel8.2 ~ 2 3 values were used for the pKa of the ligands L (DIE- corrected2* values in brackets) and the formation constants Kf of their zinc complexes : trien, pKa = 9.92, 9.25, 7.00 and 3.85 (10.53, 9.82, 7.48 and 4.28), log Kf = 12.10; 2,3,2-tet,pKa = 10.25,9.50,7.28and6.02(10.88, 10.08,7.76and6.47),logKf = 12.80; tren, pKa = 10.29, 9.84 and 8.79 (10.92, 10.45 and 9.34), logKf = 14.65. The ‘ best’ values for the acid dissociation constants of ZnL(OH2)2+ and ZnL(OD2)2+ are listed in table 1 . We estimate the uncertainty of these to be ca. & 0.1 but have not attempted a statistical error treatment for the following reasons. Although in all cases the combination of [OH-] and “a+] at the high-pH end indicates the need for a sodium-ion correction, comparison of the calculated curves with the experimental points suggests that, if anything, the fit is better with the uncorrected points than the corrected.(This is not particularly remarkable13 since the size of the alkaline error of a glass electrode varies between individual electrodes of the same type and depends on the condition of the surface. It appearsz5 to be similar in D 2 0 and H 2 0 at the same ‘measured pH’ but its dependence on the isotopic composition of the solvent has not been well documented.) Also, the shape of the computed titration curve depends significantly on the value assumed for K,; at such high ionic strengths the activity coefficient term (yH YoH/aH,O) differs markedly from unity and means that the choice of K , must be subjective.We have used interpolated values26 for K, in H 2 0 and D20 of 1.4 x mo12 dm-6, respectively. The experimental curves (fig. 1) are all consistent with the titration of a single proton. and 3.9 x l3C N.M.R. TITRATIONS The C atoms in ala and 2,3-tri are identified as in scheme 1. The measured chemical shifts 6 of the CH,, CH, and CH carbon atoms in en, gly, ala, ace and 2,3-tri at different pD are represented by the points in fig. 2. The titration curves were computed from these data by the following method. NH2-CaH,-CbH2-NH-C,H2-CdH2-C,H2-NH, Scheme 1. Identification of C atoms in L-alanine and N-(2-aminoethyl)propane- 1,3-diarnine.1488 DIE STUDY OF PK, - s a v - 4 0 - - 3 8 1 I I I 1 I 1 0 2 L 6 8 10 12 14 PD Fig. 2 (a and 6). For caption see opposite.S.P. DAGNALL, D. N. HAGUE, M. E. MCADAM AND A. D. MORETON 1489 LL 25 2 6 h E a m v 2 3 2 2 2 1 - 0 L 8 12 PD 28r 4 8 10 12 PD Fig. 2. 13C n.m.r. chemical shifts as a function of pD. (a) gly and en, (b) aia, (c) ace and (d) 2,3-tri.1490 DIE STUDY OF PKa Table 2. ‘Best’ pK, valuesa and I3C chemical shiftsb (in ppm) for ligands en 10.76, 7.95 43.98 glY 10.37, 3.14 45.85 ala 10.46 52.57 21.54 ace 5.14 24.49 2’3-tri 10.75, 10.02, 6.96 41.00 52.10 47.30 32.75 39.80 40.86 42.60 5 1.66 17.27 21.50 39.70 49.90 46.80 27.90 39.00 - 37.64 - 41.10 - - - (C, 1 (C, 1 - - - - - - 39.60 36.55 (C,) 47.10 45.45 (C,) 46.60 46.15 (C,) 27.30 24.75 (C,) 38.90 37.65 (C,) a Ca. 20 “C, I various; the estimated errors are k 0.1 (but see text). The estimated errors are kO.1 ppm.If the exchange rate for the acid dissociation HA G H+ +A- is rapid compared with the nuclear relaxation rates 7 ‘ ~ ~ and T;l, then the observed chemical shift for a particular C atom will be the average of the chemical shifts of protonated and deprotonated forms (dHA and 6,, respectively), weighted according to the fractional populations (pHA and p , ) : 6 = P H A S,, +pA 6,. (For simplicity, the charges on subscripts have been omitted and the mixture of protonated and deuteriated forms has been referred to as protonated and represented by HA.) The fractional populations are also simply related through the acid dissociation constant Ka, so if dHA, 6, and Ka are known, a plot of 6 against pD may be generated. The pKa and 13C chemical shifts derived for the individual species by means of an iterative least-squares procedure are listed in table 2; the estimated uncertainties are kO.1 and 0.1 ppm, respectively.The assignment of the low-pH form of 2,3-tri to C, and C, shown in fig. 2(d) differs from that made previo~sly,~~ although it agrees with that of Delfini et aZ.28 Our reason for changing the assignment will be discussed in a future paper, but the shapes of the two titration curves are such that the assignment has no effect on the derived pKa. DISCUSSION Five acid dissociation constants K, may be specified for the reaction at a particular temperature. First, the ‘true’ (or thermodynamic) constants KR and KO, in H 2 0 and D20 are defined as The activities ag,, az and a: of the ‘protonated’ and ‘deprotonated’ forms are those appropriate to H,O and D20, respectively, and a,, is the corrected or ‘true’ activity of the deuterium ion.Secondly, there are the corresponding ‘apparent’ (or concentration) constants K& and Kb measured at a particular ionic strength l i n termsS. P. DAGNALL, D. N. HAGUE, M. E. MCADAM AND A. D. MORETON 1491 of the concentrations of the two forms of the acid and the activity of the hydrogen (deuterium) ion : a [A-ID and K b = L [DAID 1 Kg = K R and KO, are determined by the extrapolation of K& and KL, respectively, to I = 0. Finally, it is sometimes convenient to define a 'measured apparent' constant in D,O, KE, which uses pH-meter readings (aE) directly rather than the corrected values (the correction is primarily for the junction potential between D20 and H 2 0 solutions; its magnitude depends on the concentration scale chosen25 and is usually13 0.40 added 1.0 the values measured with a pH meter standardised against ordinary pH buffers): whence pKb = p m +0.40.The ' true' and ' apparent' constants in each solvent are related to one another through the activity-coefficient ratios : and similarly Strictly, the equilibrium DIE ApK is defined as ApK = (pKO,-pKR) but it is also equal to (pKC, - pKg) if it can be assumed that the activity-coefficient ratios are equal in solutions which differ only in the nature of the solvent. This is not unreasonable since the activity coefficients of neutral molecules will be close to unity and those of ions will be very similar as the permittivities of H 2 0 and D,O are almost identical (78.39 and 78.06, respecti~ely,~~ at 25 "C).In the present discussion we shall be comparing our measured pK& and literature or measured pKC, at similar I. Although no general rationalisation of the large amount of DIE data in the literature has found acceptance, there have been several correlations of restricted samples, of which we shall mention three. Robinson et al.24 reviewed critically their data on 15 acids and concluded that ApK varied linearly with pKR at pKR > 7; two strong inorganic acids lay on an extension of this line but most of the points, representing 9 organic acids with pKR < 7, corresponded to 0.5 5 ApK 5 0.6. Ohtaki and MaedalO also noted the approximate constancy of ApK for many organic acids: considering the data for 39 compounds of four types (acetic and malonic acid derivatives, phenols and substituted anilinium ions) covering a wide range of PKH (ca. 2-6, 1-9, 3-10 and -1-4.5, respectively), they found in all cases a remarkably small spread of ApK [0.49 0.06, 0.48 0.06, 0.55 f: 0.06 and 0.56 f.0.03 (errors are standard deviations)]. Amaya et al.ll pointed to a similar result for the deprotonation of cationic metal species: although the DIE has only been reported in a dozen or so cases, ApK often seems to be ca. 0.7. In table 3 we have listed our measured pK5 for 11 acid-base equilibria, together with the measured or literature values of pKC, and the calculated ApK. The agreement with literature values (gly and ace) is good. In all cases a similar DIE is observed to those quoted;l0> 11, 24 in fact, treating our data as1492 DIE STUDY OF PK, Table 3.Comparison of measured and literature pK, values in H,O and D,O and calculated DIE species PKb pK& ApK(ca1c.) en ala ace 2.3-tri Zn(trien) (OH,)2+ Zn(2,3,2-tet) (OH2),+ Zn(tren) (OH2),+ 10.76" 7.95" 1 0.37"~ 3.14". 1 0.46" 5. 14" 10.75" 1 0.02a 6.96" 9.74" 10.91" 11.21" 10.19' 7.41 9.73d 2.41d 9.79f 4.559 10.24i 9.24i 6.28i 9.07" 10.30" 10. 59" 0.57 0.54 0.64e 0.73 0.67 0. 59h 0.51 0.78 0.68 0.67 0.61 0.62 Present study. Ref. (43) ( I = 0.65, 25 "C) and (44) ( I = 0.4-0.6, 25 "C). Ref. (45) ( I = 0.1 1, 25 "C) gives 10.28 and 2.86 and ref. (46) ( I = 0.01, 20 "C) gives 10.38 and 2.75. Values of 0.59 [ref. @)I, 0.63 [ref. (45)] and 0.53 [ref. (46)] have been reported.f Ref. (47) ( I = 1,20 "C) and (48) ( I = 0.15, 25 "C). Values of 0.46 [ref. (46)], 0.52 [ref. (51) and (52)] and 0.56 [ref. (24)] have been reported. Ref. (21) and (22) ( I = 0.5, 25 "C). Ref. (44) ( I = 0.4-0.6, 25 "C) and (47) ( I = 1 .O, 20 "C). Ref. (49) ( I = 1, 20 "C) and (50) ( I = 0.1, 20 "C). a single sample yields a mean ApK of 0.63 with a standard deviation of 0.07. In view of this, we shall in future12 apply a correction term ApK of 0.63 to the literature value of pKC, in all cases where we have no direct measurement of pKL. The effect of ionic strength on the pKa for the removal of a proton from a molecule or ion is given by log (Yp/Yd), where yp is the activity coefficient of the protonated form and yd that of the deprotonated form.For neutral molecules, log y p , d is given by kl, where k is the 'salting-out' coefficient,26 while for ions it may be estimated by eqn (1). If either species is neutral, the correction term log(y,/y,) will be reasonably small and constant, even at quite high ionic strengths. Thus, at I = 0.1 mol dm-3, log (yp/yci) is kO.11, at I = 0.5 mol dm-3 it is k0.14 and at I = 1.0 mol dm-" it is kO.10. If, however, both species are ionic (with charges zp and zd) the correction term becomes and can in principle rise rapidly with increasing ionic strength, as the example of 2,3-tri illustrates. This molecule contains three well separated amino groups which are protonated during the titration; the doubly and fully protonated ions therefore carry net charges of 2 + and 3 + , respectively, and it makes a considerable difference to the correction term log (yp/yd) whether they are treated as single, multiply charged ions or as two or three singly charged ions.On the former basis, at the equivalence point for the addition of the second proton the ionic strength of a 0.2 mol dm-3 (in 2,3-tri) solution will be 0.4 mol dm-3, giving a pKa correction term of -0.40 (assuming26 A = 0.51); at the equivalence point for the addition of the third proton the ionic strength will be 0.9 mol dm-3 and the correction term -0.62. If, on the other hand, they are treated as 'hybrid' ions13 the ionic strengths become 0.3 and 0.5 mol dm-3,S . P. DAGNALL, D. N. HAGUE, M. E. MCADAM AND A. D. MORETON 1493 respectively, and the value of log (yp/yd) is -0.14 in each case.The values of ApK for these two protonations (0.78 and 0.68, respectively) therefore provide evidence in favour of treating these multiply charged ions as separate unit charges, since the values of pKC, were determined in solutions of I = 0.5 mol dm-, KNO, [ref. (22)] or KCl [ref. (21)] where no such ambiguity exists. The coordination geometries in the zinc complexes of the tetramines trien, 2,3,2-tet and tren in aqueous solution have been the subject of debate [see, for example, ref. (21) and (30)-(36)] but it seems to be assumed, largely on thermodynamic grounds, that the metal adopts a coordination number of four in at least one case. We have already provided3' X-ray diffraction evidence that the trigonal-bipyramidal configuration e s t a b l i ~ h e d ~ ~ - ~ ~ for Zn(tren) in the solid state is retained in solution but no direct evidence has been published for the other two.[Solid zinc(trien) iodide has been shown4' to contain square-pyramidal Zn(trien)I+ ions.] The observance of a normal DIE (table 3) for all three complexes indicates that the [ZnL];: ions all contain an acidic water molecule and therefore that the coordination number of the metal is at least 5 : if the hydroxide ion were adding to the zinc complex rather than removing a proton from it, the effect of isotopic substitution would be much smaller42 since the reaction would not then involve the rupture of a bond to hydrogen. We thank I.C.I. for a Research Fellowship (to M. E. M.) and the S.R.C./S.E.R.C. for studentships (to S.P.D.and A. D. M.). ' C. W. Davies, Electrochemistry (Newnes, London, 1967), p. 60. C. W. Davies, in The Structure of Electrolyte Solutions, ed. W. J. Hamer (Wiley, New York, 1959), p. 23. P. M. Laughton and R. E. Robertson, in Solute-Solvent Interactions, ed. J. F. Coetzee and C. D. Ritchie (Dekker, New York, 1969), p. 399. C. K. Rule and V. K. La Mer. J. Am. Chem. Soc., 1938, 60, 1974. R. P. Bell, The Proton in Chemistry (Methuen, London, 1959, p. 188). A. 0. McDougall and F. A. Long, J. Phys. Chem., 1962, 66, 429. ' R. P. Bell and A. T. Kuhn, Trans. Faraday Soc., 1963, 59, 1789. W. P. Jencks and K . Salveson, J. Am. Chem. Soc., 1971,93, 4433. E. Hogfeldt and J. Bigeleisen, J. Am. Chem. Soc., 1960, 82, 15. lo H. Ohtaki and M. Maeda, Z . Naturforsch. Teil B, 1972, 27, 571.l 1 T. Amaya, H. Kakihana and M. Maeda, Bull. Chem. Soc. Jpn, 1973, 46, 2889. l 2 S. P. Dagnall, D. N. Hague, M. E. McAdam and A. D. Moreton, to be published. l 3 R. G. Bates, Determination of p H , Theory and Practice (Wiley, New York, 1964). l 4 L. J . Wilson and N. J. Rose, J. Am. Chem. Soc., 1968,90, 6041. A. I. Vogel, Quantitative Inorganic Analysis (Longmans, London, 3rd edn, 1961), p. 433. R. M. C. Dawson, D. C. Elliott, W. H. Elliott and K. M. Jones, Data for Biochemical Research (Clarendon Press, Oxford, 2nd edn, 1969), chap. 20. l 7 A. Albert and E. P. Serjeant, Ionization Constants of Acids and Bases (Methuen, London, 1962), chap. 2. W. A. C. McBryde and H. K. J. Powell, Can. J . Chem., 1979, 57, 1785. l 9 M. T. Beck, Chemistry of Complex Equilibria (van Nostrand, London, 1970), chap.4. 20 L. G. Sillen and A. E. Martell, Stability Constants of Metal-ion Complexes (The Chemical Society, London, 1964); Stability Constants of Metal-ion Complexes, Supplement 1 (The Chemical Society, London, 1971). 21 D. C . Weatherburn, E. J. Billo, J. P. Jones and D. W. Margerum, Znorg. Chem., 1970, 9, 1557. 2 2 R. Barbucci, L. Fabbrizzi and P. Paoletti, Inorg. Chim. Acta, 1973, 7, 157. 23 G. Anderegg and P. Blauenstein, Helv. Chim. Acta, 1982, 65, 913. 24 R. A. Robinson, M. Paabo and R. G. Bates, J . Res. Natl Bur. Stand., Sect. A , 1969, 73, 299. 25 A. K. Covington, M. Paabo, R. A. Robinson and R. G. Bates, Anal. Chem., 1968,40, 700. 26 H. S. Harned and B. B. Owen, The Physical Chemistry of Electrolytic Solutions (Reinhold, New York, 27 S.P. Dagnall, D. N. Hague and M. E. McAdam, J. Chem. Soc., Perkin Trans. 2, 1984, 1 1 11. 3rd edn, 1958).1494 DIE STUDY OF PK, 28 M. Delfini, A. L. Segre, F. Conti, R. Barbucci, V. Barone and P. Ferrute, J . Chem. SOC., Perkin *$ G. A. Vidulich, D. F. Evans and R. L. Kay, J . Phys. Chem., 1967,71, 656. 30 H. Ackerman, J. Prue and G. Schwarzenbach, Nature (London), 1949, 163, 723. 31 H. Irving and R. J. P. Williams, J. Chem. Soc., 1953, 3192. 32 L. Sacconi, P. Paoletti and M. Ciampolini, J. Chem. Soc., 1961, 51 15. 33 P. Paoletti, M. Ciampolini and L. Sacconi, J. Chem. Soc., 1963, 3589. 34 R. Barbucci, L. Fabbrizzi, P. Paoletti and A. Vacca, J. Chem. SOC., Dalton Trans., 1973, 1763. 35 L. Fabbrizzi, R. Barbucci and P. Paoletti, J . Chem. SOC., Dalton Trans., 1973, 1529. 36 P. Paoletti, L. Fabbrizzi and R. Barbucci, Znorg. Chim. Acta Rev., 1973, 7 , 43. 37 S. P. Dagnall, D. N. Hague and A. D. C. Towl, J . Chem. SOC., Faraday Trans. 2, 1983, 79, 1817. 38 G. D. Andreetti, P. C. Jain and E. C. Lingafelter, J . Am. Chem. SOC., 1969, 91, 41 12. 39 R. J. Sime, R. L. Dodge, A. Zalkin and D. H. Templeton, Inorg. Chem., 1971, 10, 537. 40 M. Duggan, N. Ray, B. Hathaway, G. Tomlinson, P. Brint and K. Pelin, J. Chem. Soc., Dalton 41 G. Marongiu, M. Cannas and G. Carta, J . Coord. Chem., 1973, 2, 167. 42 E. J. King, Acid-Base Equilibria (Pergamon, Oxford, 1965), p. 265. 43 F. Basolo and R. K. Murmann, J. Am. Chem. SOC., 1952,74, 5243. 44 D. L. Rabenstein and G. Blakney, Inorg. Chem., 1973, 12, 128. 45 N. C. Li, P. Tang and R. Mathur, J. Phys. Chem., 1961, 65, 1074. 46 G. Schwarzenbach, A. Epprecht and H. Erlenmeyer, Helv. Chim. Acta, 1936, 19, 1292. 47 D. D. Perrin, J . Chem. Soc., 1958, 3120. 48 C. Tanford and W. S. Shore, J . Am. Chem. SOC., 1953,75, 816. 49 D. D. Perrin, J . Chem. SOC., 1959, 1710. 50 R. S. Kolat and J. E. Powell, Inorg. Chem., 1962, 1, 293. 5 l V. K. La Mer and J. P. Chittum, J. Am. Chem. Soc., 1936, 58, 1642. 52 P. K. Glasoe and F. A. Long, J . Phys. Chem., 1960, 64, 188. Trans. 2, 1980, 900. Trans., 1980, 1342. (PAPER 4/ 1733)

ISSN:0300-9599    

DOI:10.1039/F19858101483

出版商:RSC    

年代:1985

数据来源: RSC

摘要:

J. Chem. SOC., Faraday Trans. I , 1985,81, 1495-1501 Partial Molar Isobaric Heat Capacities of a Substance in a System containing a Chemical Equilibrium A Basis for the Estimation of the Effect of Solvent Reorganisation on Heat Capacities of Activation for Chemical Reaction BY MICHAEL J. BLANDAMER* AND JOHN BURGESS Department of Chemistry, University of Leicester, University Road, Leicester LE 1 7RH Received 1st October, 1984 The partial molar enthalpy and partial molar isobaric heat capacity of a substance Z is considered for a system containing a chemical equilibrium between substances X and Y. An equation for the isobaric heat capacity of Z identifies two contributions. The first describes the heat capacity of Z when the chemical reaction between X and Y is frozen whereas the second describes a contribution from a change in composition to preserve chemical equilibrium between X and Y.The treatment offers a method for quantifying the effects of solvent reorganisation on the heat capacities of activation for chemical reaction. The isobaric heat capacity of a closed system containing a chemical equilibrium can show a striking dependence on temperature. This dependence arises because the position of chemical equilibrium depends on temperature and shifts when heat is absorbed.' Recently2 we considered the case where a system contains two coupled equilibria, e.g. X e Y e Z where X, Y and Z are three chemical substances. Here we consider an intermediate case where an inert substance (e.g. Z) is present in a system in which two substances X and Y are in chemical equilibrium.We explore the impact of the chemical equilibrium X e Y on the partial molar isobaric heat capacity of substance Z. Previously2 we assumed that, other than the chemical equilibria involving X, Y and Z, these three substances are ideal in a thermodynamic sense. If this same assumption is used here, the molar properties of Z are insensitive to the equilibrium between X and Y. Some element of non-ideal behaviour must be incorporated into the analysis in order that substances X and Y communicate with Z. Originally we had in mind a mixture of gases with the intention of extending the analysis toward our interests in understanding heat capacities of activation for reactions in solution. Explanations of the sign and magnitude of A:Cg have aroused controversy, particularly for reactions in water and aqueous mixtures. We therefore envisaged a solvent formed from two liquids (e.g.water and ethyl alcohol) which form a non-ideal mixture. However, we have modified our approach in the light of a recent important paper by Gr~nwald,~ who uses the equilibrium between X and Y as a two-state model for liquid water, substance Z being a solute. The treatment given here differs from that offered by Gr~nwald,~ who uses a lattice-type model for water. 14951496 ISOBARIC HEAT CAPACITIES DESCRIPTION OF SYSTEMS A vital consideration is the interplay between the description of a system and characterisation of substances within the system. We use two descriptions of the system under discussion (at temperature T and pressure p ) .In description I the system contains n, moles of solvent S and n, moles of solute Z. The extensive Gibbs function of state G is defined by the independent variables T, p , n, and n,: (1) Eqn (1) can be rewritten using, for example, the volume V and the enthalpy H rather than G. Further, G = G[T; p ; n,; n,]. .=[El T ; n s ; n, (2) In description I1 the solvent S is a mixture of two substances X and Y in chemical (3) ,uEq (system; T ; p ) = ,u$q (system; T ; p ) (4) and n, = nzq +n;q. ( 5 ) equilibrium : where X+Y The equilibrium composition of the solvent is characterised by the extent of chemical reaction ceq such that n$q = l e q and ngq = n , - t e q . Here <, an extensive property, measures the composition of the solvent relative to that for a liquid comprising n, moles of pure X.The Gibbs function is given by G = G[T; p ; ngq; n$q; n,] (6) or G = G[T; p ; n, ; t e q ; n,]. (7) Eqn (l), (6) and (7) are formulations of the Gibbs function for the system. No matter how we, as observers, describe a given system all descriptions must be consistent with the same macroscopic extensive function of ~ t a t e . ~ (A similar state of affairs occurred when the blind men of Hindustan encountered an elephant.) PARTIAL MOLAR PROPERTIES The partial molar properties of Z are probed by examining how the macroscopic properties of the system change when dn, moles of Z are added. For example, the partial molar volume of Z, V,, is measured as the resulting change in volume 6V at equilibrium for the systems having volumes V and V+dV: In eqn (8) we used description I.However complicated our description of the equilibrium system, the quantity V, is description-independent. Nevertheless descrip- tion I1 raises the possibility of different interpretations of V,. We define two new quantities using eqn (9) and (10): T ; p ; nxeq ; nyeq (9) T ; p ; n , ; A - OM. J. BLANDAMER AND J. BURGESS 1497 Eqn (9) and (10) carry additional information characterising the chemical reaction between X and Y. Either the amounts of X and Y do not change, remaining at ceq [eqn (9)], or change in order to hold the system at equilibrium, where the affinity for spontaneous change, A , is zero [eqn (lo)]. In analogous fashion, partial molar isobaric heat capacities of Z can be defined at either Ceq or ' A = 0' ; e.g.eqn (1 1) for C,,(A = 0), [cf. eqn. (lo)]: (1 1) Faced with the problem of interpreting measured V, and Cpz, we may envisage using description I1 for two contributions to a given partial molar property. The first contribution measures, for example, the change in C, in the event that the amounts of X and Y do not change. This contribution is called Cpz(c), the 'frozen' or 'instantaneous' partial molar property. The difference between C,,(t) and the equilibrium partial molar isobaric heat capacity C,,(A = 0), defined by eqn (1 l), represents the relaxation of the solvent to the new equilibrium state. In the analysis presented below we formulate an expression for C,,(A = 0) in terms of these two contributions. Implicit in the analysis is the assumption that experiment reports an equilibrium property.SOLVENT EQUILIBRIUM The solvent is described (description 11) as a mixture of two liquids X and Y. The chemical potential of either X or Y is related to the mole-fraction composition using eqn (12) with i = X or Y: pi(system; T;p) = p; (1; T)+RTln ( x i f i ) + Vc (1; T;p) dp where, by definition, lim (xi -+ 1.0) fi = 1 at all T and p. The standard substance i is the pure liquid at temperature T and pressure p". The chemical of solute Z is related to the molality m, using the equation L f P (12) state for potential p,(system; T ; p ) = pi (sln; T)+RTln (m, y,/m")+ VF (sh; T;p) dp (13) where lim (m, + 0) yz = 1.0 at all T and p ; mo = 1 mol kg-l. The standard state for Z is a solution in a solvent comprising an equilibrium mixture of X and Y where m, = 1 and yz = 1 at temperature Tand pressurep".The assumption is made that the activity coefficients for the liquid components are related to the molality of Z in solution. We adopt the procedure3 where both lnfx and lnf, are linear functions of m,: J P O lnfx = 8xmz/m0 (14) w, = 8, mz/mo (15) and In (f,/fx) = Pmz/m"* (16) However, we assume that Z in the solvent mixture is an ideal solute at all T and p , i.e. yz = 1. Hence derived properties for solute Z in solution where m, = 1 are the corresponding molar properties; e.g. the molar volume of Z is V,cC, (sln; T ; p ) .1498 ISOBARIC HEAT CAPACITIES GIBBS FUNCTION AND CHEMICAL POTENTIAL The chemical potential of substance Z measures the change in G, dG, when dn, moles of Z are added to the system.We envisage two limiting conditions under which (i) 5 and (ii) the affinity for spontaneous change, A , for the solvent reaction are constant. The partial differentials of G with respect to n, under these two constraints are related: where The triple-product term on the right-hand side of eqn (1 7) is of immediate interest. For a system exhibiting bilateral stability, (aA/at) < 0 in the region near chemical equilibrium. If the constraint is imposed on eqn (17) that the chemical reaction between the solvent components is at thermodynamic equilibrium (i.e. A = 0), this triple-product term disappears. Hence, with reference to the chemical potential of 2, = [El'" . [:IT; p ; n,; A-0 T ; P ; n,; t In other words the chemical potential of Z in the system is insensitive to a description of the solvent as comprising either one substance S or two substances X and Y in equilibrium.In effect the condition given by eqn (4) feeds through into the definition of the chemical potential for Z [cf. eqn (1 3)]. Eqn (17) can be rewritten in terms of the partial derivatives of G with respect to the variables T and p . For example, [using eqn (1 8)] For a system at equilibrium (where A = 0) the volume is not a function of how the system is described. Similar comments apply to the enthalpy H, which is also related to a first derivative of the Gibbs function. ENTHALPIES AND HEAT CAPACITIES Although the chemical potentials of X and Y at equilibrium are equal [eqn (4)] the corresponding partial molar enthalpies, entropies and heat capacities are not neces- sarily equal. In all probability they are quite different.If eqn (1 7) is written as a partial derivative of the volume or enthalpy, the corresponding triple-product terms are not zero at equilibrium: [ E I T ; p ; n,; AM. J. BLANDAMER AND J. BURGESS 1499 2V/a< and aH/a< at constant T;p;n,;n, and at equilibrium are not zero. To understand the significance of the latter statement it is informative to calculate the three quantities in the triple product term in eqn (22). The affinity for spontaneous change within the solvent reaction is given by A = - [py(system; T;p)-px(system; T;p)]. xy = ny/(ns+n,) = l/%. (23) (24) For a solution which is dilute in Z (with dnj = vjdt, where vj is the stoichiometry) Similarly xx = ( n s - < ) / ~ s .(25) Hence using eqn (12) and (23)-(25) = -RT[n,/(n,-<) c] = RT/n,(l -xy) xy (26) or In deriving eqn (27) we used eqn (16). Further, For a solution dilute in Z nzz = nz/(nx M x + ny My) where Mx and My are the molar masses of substances X and Y. Then where Vk' is the mass of solvent. For a dilute solution of Z and at fixed n, (assuming PX and PY are independent of temperature) = (G -%) = A, H* (31) [ g ] T ; p ; n , ; n , where HE and G are the molar enthalpies of pure X and Y at temperature T and pressure p. Combination of eqn (22), (27), (28) and (31) yields the partial molar enthalpy of Z in a system where X and Y are in equilibrium: Hz(A = 0) = HZ(leq) - (n, P/mo W ) x;q( 1 - x;q) A, H*. (32) The differential of eqn (32) with respect to temperature yields the (equilibrium) partial molar isobaric heat capacity of substance Z.If 4 = n, P/mo W C p , ( A = 0 ) = Cp,(<eq) - 4(d/dT) x:' (1 - x$q) Ar h*. (33) (34) then, at constant n,, n,, T and p,1500 Hence, ISOBARIC HEAT CAPACITIES C,,(A = 0) = C,,(ceq)-4x$q (1 -x$q)Ar Cg-gArH* (d/dT) [x$q(l -x$')]- (35) At a given T and p , the first two terms on the right-hand side of eqn (35) reflect the contributions of C,,(leq) and the solvent reaction to C,,(A = 0). The third term (by definition, Cp",) registers the sensitivity of the solvent reaction to a change in temperature : (36) Cgz = -$A, H* (1 -2 ~ $ 9 ) (dx$q/dT)p. The molar enthalpies of X and Y are assumed to be independent of temperature. The algebra is simplified by assuming that ambient pressure p equals the standard pressure PO.Therefore, A, H" = Wy - W, = A,. H*. The equilibrium constant for the solvent reaction K"( T ) = p / ( n , - p q ) (37) where A, Go( T ) = - RT In K"( T ) = P i (1 ; T)-&l; T ) (38) and at p o A,H" = Wy(l; T)-Wx(l; T ) . (39) dceq/d T = (n, K"( T ) / [ 1 + K"( T)I2) A, F / R T 2 . (40) Hence c:, = -{@(I -2x;q)K"(T)/[l + P ( T ) l 2 ) [A,H0I2/RT2- (41) Also [eqn (37)] K"( T ) / [ 1 +A?( T)]2 = x;q (1 - xyl). (42) Cg, = - (n,/3/mo W ) xtq( 1 - x;q) (1 - 2 x5q) [A, H"I2/RT2. (43) Then Then, in a form analogous to that obtained by Gr~nwald,~ Here (n,/W) is the number of moles in 1 kg of solvent and, perhaps significantly, is larger for water than for organic solvents. The overall forms of eqn (41) and (43) hide a complex dependence on temperature.Even assuming ArW is independent of temperature, C$, is dependent on temperature because ceq depends 011 temperature [cf. eqn (40)]. These recursive features in the equations for derived C , quhrtities have been commented on previously.2 DISCUSSION There are marked similarities between the equations developed here for the isobaric partial molar heat capacity of 2, C,,(A = 0), and those reported previously2 for the molar heat capacity of a system containing an equilibrium. This is not unexpected in the sense that here substance Z probes the changes occurring in the system by virtue of the solvent reaction. In eqn (41) the term K"/( 1 + K"), together with the square of A, H", is a direct consequence of, in Hammett's term^,^ the adoption of a sophisticated model for a system.The quest for this overall pattern, particularly with respect to eqn (41), prompted the analysis reported here. If a different assumption is invoked for the dependence off, andf, on m, [cf. eqn (14)-(16)], the final equation for C,,(A = 0) differs in detail but not in substance. In the context of aqueous solutions5 /3 measures the extent to which Z stabilises or destabilises the solvent components X [ = (H,O),, the bulky hydrogen-bonded state] and Y [ = (H20),, the dense non-hydrogen-bonded state], i.e. structure breaking orM. J. BLANDAMER AND J. BURGESS 1501 structure making. A shift in the solvent equilibrium cannot account for trends in chemical potentials. However, trends in C,, ( A = 0) and H , ( A = 0) may be understood in these terms.The treatment can be modified for a description of the properties of solutes using the more sophisticated two-state model for water discussed recently by Lumry et aL6 [see also heat-capacity data in ref. (7)]. The analysis may also be extended to related systems. Substances X and Y may describe two liquid cosolvents where a solute is two substances Z and Z' in equilibrium; e.g. the active and denatured forms of a protein or two conformers of a carbohydrate in aqueous solution. Further, Z and Z' may describe two solvates Z - h X and Z(h- 1)X.Y in equilibrium, linking the analysis with preferential solvation models developed by Covington et aZ.** and applied recently by Remerie and Engberts.lO With regard to our own interests, attention is directed towards the application of these equations to kinetic data and the interpretation of the isobaric heat capacity of activation AIC:, where solute Z models first the initial and then, as Zi, the transition states.Eqn (35) and (41) provide the first method, as far as we are aware, of describing quantitatively solvent effects on AS H a and AICg in terms of a coupled solvent reaction. Interestingly the final form of the equations is similar to that obtained using the Albery-Robinsonll mechanism for the solvolysis of t-butyl chloride in water where instead of K"(T) we write a, the ratio of two rate constants. The obvious difference is the starting hypothesis. The Albery-Robinson scheme expresses the overall reaction as proceeding through an intermediate.Although thermodynamics cannot resolve the controversy concerning the source of the negative A'CP value for this class of reactions,12 we are currently considering how the equations described above may be applied to particular systems. Analogous arguments may be advanced in connection with the dependence of Ar H a on T for the dissociation of weak acids in aqueous s01utions.l~ E. D. McCollum, J. Am. Chem. SOC., 1927,49, 28. * M. J. Blandamer, J. Burgess and J. M. W. Scott, J. Chem. SOC., Faraday Trans. I , 1984, 80, 2881. E. Grunwald, J. Am. Chem. SOC., 1984, 106, 5414. L. P. Hammett, Physical Organic Chemistry, (McGraw-Hill, New Yotk, 2nd edn, 1970), chap. 2. F. Franks, Water--A Comprehensive Treatise (Plenum Press, New York, 1973), vol. 1, chap. 1 . R. Lumry, E. Battistel and C. Jolicouer, Faraduy Discuss. Chem. SOC., 1982, 17, 93. ' M. Oguni and C. A. AngelI, J. Chem. Phys., 1980,73, 1948. * A. K. Covington, T. H. Lilley, K. E. Newman and G. E. Porthouse, J. Chem. SOC., Faraday Trans. 1 , 1973, 69, 963, A. K. Covington, T. H. Lilley and K. E. Newman, J . Chem. SOC., Faraday Trans. I , 1973,69, 973. lo K. Remerie and J. B. F. N. Engberts, J. Phys. Chem., 1983,87, 5449. I 1 W. J. Albery and B. H. Robinson, Trans. Faruday SOC., 1969, 65, 980. ** M. J. Blandamer, J. Burgess, P. P. Duce, R. E. Robertson and J. M. W. Scott, J. Chem. SOC., Faraday l 3 M. J. Blandamer, J. Burgess, R. E. Robertson and J. M. W. Scott, Chem. Rev., 1982, 69, 259. Trans. 1, 1981, 65, 1999. (PAPER 4/ 1697)

ISSN:0300-9599    

DOI:10.1039/F19858101495

出版商:RSC    

年代:1985

数据来源: RSC

摘要:

J. Chem. Soc., Faraday Trans. I, 1985, 81, 1503-1512 Thermodynamics of Adsorption of 0-n-Dodecylpentaethylene Glycol and 0-n-Dodecyloctaethylene Glycol from Aqueous Solutions on to Graphitised Carbon BY ANDREA GELLAN AND COLIN H. ROCHESTER* Department of Chemistry, The University, Dundee DD1 4HN Received 5th October, 1984 Adsorption isotherms and enthalpies of adsorption are reported for two dodecylpolyethylene glycol non-ionic surfactants (C,,E, and C,,E,) from aqueous solution on to graphitised carbon. Adsorption at low coverages involved both the hydrophobic and hydrophilic parts of the surfactant molecules lying flat on the carbon surface. Subsequent displacement of adsorbed oxyethylene segments by incoming dodecyl groups was accompanied by dehydration of the hydrophilic groups.Segments of adsorbed alkyl groups were lifted from the surface as the adsorbed layer approached the maximum attainable surface excess concentration. At this stage the differential enthalpy of adsorption of C,,E, on carbon was similar to a literature value for the standard enthalpy of adsorption of C,,E, at the water/air interface. Isotherms for the adsorption of 0-alkylpolyethylene glycols from aqueous solution on to the surface of solids are influenced by the length of the hydrocarbon chains and the number of ethylene oxide segments in the non-ionic surfactant molecules.' Models of adsorption invoke surfactant molecules lying parallel or perpendicular to the adsorbent surface, the latter becoming more significant with increasing surface coverage. At intermediate stages of adsorption either the polar or non-polar groups in the adsorbed surfactant may be preferentially displaced from the surface as a result of adsorbate-adsorbate interactions in the surface layer.High equilibrium surfactant concentrations in solution may result in the formation of absorbed bilayers or, by analogy to micellisation in solution, adsorbed micelles or half-micelles. Models of adsorption have primarily been developed from free-energy (isotherm) data. Enthalpy and entropy data might, however, be more sensitive to changes in adsorption behaviour because the free energies will result in part from compensating contributions from the enthalpy and entropy changes accompanying adsorption. As only few calorimetric studies of adsorption of non-ionic surfactants have been r e p ~ r t e d , ~ ~ ~ we have undertaken a study of the adsorption of three 0-alkylpolyethylene glycols from aqueous solutions on to model hydrophobic (graphitised carbon) and hydrophilic (silica) surfaces.Corkill et al.3 reported the enthalpies of adsorption on Graphon of a series of 0-n-alkylhexaethylene glycols (C,E,) with alkyl chain lengths (x) of 6, 8, 10 and 12. The present paper reports calorimetric results for the immersion of graphitised carbon in aqueous solutions of two surfactants with the same alkyl chain length (C12) but with either five (C12E5) or eight (C,,E,) ethylene oxide segments constituting the polar part of the surfactant. Our results for the adsorption of 0-n-octyltetraethylene glycol (C,E,) on graphitised carbon have already been published.6 15031504 ADSORPTION OF SURFACTANTS EXPERIMENTAL 0-n-Dodecyloctaethylene glycol (C12E,) and 0-n-dodecylpentaethylene glycol (C12E,) were used as supplied by the Nikko Chemical Co., Tokyo, and were storedprior to use under nitrogen at ca.0 "C. Vulcan-3-G (2700) graphitised carbon black was used as provided by the National Physical Laboratory as standard-surface-area material with an area (nitrogen B.E.T.) of 71.3 & 2.7 m2 g-l. Quoted levels of impurity were ash 0.06%, Fe203 0.005%, Ca 0.0025%, Na 0.002%, K 0.005%, S 0.009% and traces (< 0.0005%) of Mg, Si, V and Al. Transmission electron micrographs taken with a JEOL JEM200CX instrument showed the carbon to be fully graphitised with an approximate diameter of polyhedral primary particles of 20-40 nm and an interlayer spacing of 0.34 nm.Lattice-plane defects were also observed in the high-resolution micrographs. Water was deionised and then triply distilled under nitrogen, once from alkaline potassium permanganate and twice from itself. Adsorption equilibria were established by rotation of ca. 0.2 g aliquots of carbon and 10 cm3 of aqueous surfactant solution in stoppered glass tubes mounted on an end-over-end stirrer immersed in a water thermostat at 25.00+0.05 "C. Separation of solid and solution was achieved by settling or centrifugation. Solution removed by pipette was diluted, when necessary, to a concentration below the surfactant c.m.c. and estimated by measurement of surface tension using the detachment Wilhelmy plate technique.Three microscope cover slips (cross-section dimensions 32 x 0.15 mm) were suspended in solutions from one arm of a modified Stanton model B16 balance and the weight on the second balance arm was increased (minimum increment 0.1 mg) until detachment occurred. The balance was housed in an air thermostat maintained at 25.0k0.1 "C. The measured surface tension of purified water was 71.93 f0.06 mN m-' (literature' value 71.90 mN m-l). Calibration graphs of surface tension against log (concentration) gave c.m.c. values of 6.53 x lo-, mol dmP3 for C,,E, and 1.05 x loP4 mol dmP3 for Cl,E, at 25 "C. Enthalpies of immersion of graphitised carbon in aqueous surfactant solutions and of dilution of concentrated solutions (0.1 mol dm-3 Cl,E, and 0.5 mol dm-3 C,,E,) with 100 cm3 of water were determined with an LKB 8700-2 reaction calorimeter, as before,6 with the thermostat bath maintained at 25.00 f 0.05 "C.RESULTS Values of critical micelle concentration (c.m.c.), Gibbs free energy of micellisation (AG&c), surface tension at the c.m.c. (ycmc), maximum surface excess concentration (r) and minimum area per molecule ( A ) for C12E5 and C12E, in water at 298 K were deduceds from plots (fig. 1) of surface tension against log,, (concentration) and are compared in table 1 with existing literature data. The AGkic values were calculated as RTln (c.m.c.) with c.m.c. on the mol dm-3 concentration scale. Isotherms for the adsorption of CI2E5 and C&, on Vulcan-3-G at 298 K are shown in fig. 2. The isotherm for C12E, resembles that previously reported6 for C,E4 in that there is no point of inflection and a plateau is reached at the c.m.c.Similarly shaped isotherms have been obtained for several 0-n-alkylpolyethylene glycols adsorbed from aqueous solution on to G r a p h ~ n . ~ ? l2 The plateau for C&8 corresponds to a surface concentration of 2.16 x loplo mol cmp2, which is equivalent to an average surface area per molecule of 0.77 nm2, approximately twice the value of 0.39 nm2 for C8E4.6 The present results for C12Es are closely consistent with the isotherm of Lange12 for C,,E, on Graphon. At low equilibrium concentrations of surfactant the isotherm for C,,E5 increased less steeply than that for C12E,. However, the onset of an adsorption plateau for C12Es was not accompanied by a similar plateau for C12E5 in the same concentration range.The extent of adsorption of C12E5 continued to increase and became greater than for C12E, above a concentration of ca. 6.6 x rnol dm-3. TheA. GELLAN AND C. H. ROCHESTER I I I I 1 1505 3 0 - I I I I 1 I -7 -6 -5 -4 -3 -2 log,, (concentration/mol dm-3) Fig. 1. Variation of surface tension with surfactant concentration for aqueous solutions of (a) Cl2E5 and (6) C12E8 at 298 K Table 1. Parameters deduced from the surface tension data for aqueous solutions of Cl,E5 and C,,E, at 298 K c.m.c. AGnic Ycmc r A surfactant / 1 OW5 mol dm-3 /kJ mol-l /mN mP2 / 1 0-lo mol cmF2 / nm2 'lZE5" 6.5 - 23.9 30.4 2.98 56 C12E5* 6.4 - 23.9 30.4 3.3 1 50 C12E5e 5.6 - 24.3 31.4 3.5 48 C12E8" 9.2 - 23.0 33.9 2.43 68 C I A b 10.9 - 22.6 34.7 2.52 66 C12Ead 7.1 - 23.7 34.3 2.72 61 a Present work; ref.(8); ref. (9); ref. (10) and (11). surface concentration and the average surface area occupied per molecule for both surfactants at the cross-over point were 2.13 x 1O-lo mol cmP2 and 0.78 nm2, respec- tively. An isotherm for higher, but overlapping, concentrations of C12E5 in equilibrium with carbon black at 298 K exhibited a level of adsorption in the overlapping concentration range closely similar to the present data.13 A plateau in the adsorption of C1ZE5 on carbon black occurred in the equilibrium concentration range (6-15) x mol dm-3 and corresponded to a surface concentration of ca. 3.1 x 10-lo mol cm-2 and an average surface area per adsorbed molecule of 0.54 nm2.13 Integral heats of adsorption Qads from aqueous solutions of C,,E5 and C,,E, onI506 ADSORPTION OF SURFACTANTS 1 2 3 equilibrium concentration/ mol dm-3 Fig.2. Adsorption isotherms on Vulcan-3-G at 298 K for (a) C,,E, and (6) C,,E,. to Vulcan-3-G were evaluated, as before,6 from the measured heats of immersion Qimm via the equation Qads = Q i m r n - 4 A E f - A R i ) (1) where It is the number of moles of surfactant remaining in solution after adsorption and AHf and AHi are the molar heats of dilution of a concentrated solution of surfactant to the final and initial surfactant concentrations, respectively, in the adsorption experiments. The correction n(dgf - AHi) was negligible when the initial concentration before adsorption was below the c.m.c. but became significant above the c.m.c.when dilution during adsorption was accompanied by demicellisation.6 The variations of Qads - Qg, where Qg is the heat of immersion of Vulcan-3-G in water, with equilibrium surfactant concentration in solution are shown in fig. 3. Information about variations in the differential heat of adsorption with increasing levels of adsorption were deduced from plots of the integral heats against the amounts of surfactant adsorbed. The graphs for both surfactants (fig. 4) were consistent with highly exothermic differential heats of adsorption at very low coverages up to ca. mol crnp2. At higher surface coverages the differential heats remained exothermic but were considerably less negative. The differential heat of adsorption of C12E5 was constant within the range of coverages (0.27-2.45) x 10-lo mol cmP2, the upper limit corresponding to an equilibrium solution concentration above the c.m.c.(table 1) for which the surface concentration of adsorbed C12E, was 2.13 x 10-lo mol cm-2. In contrast, the curve for C12E, [fig. 4(6)] showed evidence for a further change in the differential heat of adsorption at ca. 1.57 x 10-lo mol cm-2 coverage, which corresponds to a solution concentration below the c.m.c. At the highest coverages studied the differential heat of adsorption was mol g-l, which corresponds to ca.A. GELLAN AND C. H. ROCHESTER 1507 I I 5 10 equilibrium concentration/ 1 0-5 mol dm-3 Fig. 3. Heats of adsorption of (a) C,,E, and (b) C,,E, on Vulcan-3-G as a function of the equilibrium surfactant concentration in solution.$2 i I 1 ,[ cb Fig. 4. Heats of adsorption of (a) C,,E, and (b) C,,E, on Vulcan-3-G as a function of the amount of surfac tan t adsorbed.1508 ADSORPTION OF SURFACTANTS Table 2. Differential enthalpies of adsorption for C,,E, and C,,ER on Vulcan-3-G at 298 K bulk conc. r A A H AHLc AH%jA surfactant mol dm-3 /10-lo mol cm-2 /nm2 /kJ mol-l /kJ mol-l /kJ mol-l 1.3- 10.3 0.27-2.45 6.2-0.68 - 9.2 9.9 0 . 6 '12E* 0.1-1.4 0.34-1.57 4.9-1.1 - 1.7 13., 3-3 1.410.0 1.57-2,13 1.1-0.78 ca. t 2 a Enthalpies of micellization and adsorption at the water/air interface from ref. (8). positive, a result which resembles the behaviour previously reported for CsE, adsorbed on Sterling FTD-4.6 The differential heat of adsorption data are summarised in table 2.DISCUSSION The presently determined surface areas occupied per molecule by C12E5 and C12E, at the aqueous/air interface at surfactant concentrations just below the c.m.c. agreed reasonably with previously published areas,*-ll were closely consistent with a plot by Lange12 of area against the number of ethyleneoxy units for 0-n-dodecylpolyethylene glycols and may be compared with similar areas of 0.56 and 0.66 nm2 for C1oE5 and C10E8, respecti~e1y.l~ The data are compatible with the general conclusions* 12* 15* l6 that increasing the ethyleneoxy chain length increases the minimum area occupied by CnEm surfactants at the water/air interface. However, Teo et all7 have recently proposed that the alkyl chain length (n) has an important influence on the minimum molecular surface area for surfactants with rn = 27 and n = 12-23.Taking the dimensions of a hydrocarbon chain as 0.5 nm and 0.127(n- 1) nm for the diameter and length, respectively, suggested that the alkyl chains were lying flat at the water/air interface. The projected area of 0.70 nm2 for n = 12 was consistent with the experimental result for C12E2, and also agrees with the present result for C$s. Despite this agreement the near identity of the areas for C12E27 and C12E8 is surprising by comparison with results for C,E, with rn = 4-30 and n = 1218 and with m = 6-21 and n = 16,19 which showed significant increases in area with increasing rn at constant n. Plots of area against rn for the C12Em and C16Em surfactants suggest, in accordance with the results of Teo et al.,17 that increasing the alkyl chain length from n = 12-16 caused a significant increase in minimum area of surface occupied per molecule for high values of rn.However, the areas quoted by Teo et aI.l7 were lower than would be consistent with values for rn = 27 interpolated from the plots of area against rn for the earlier data. In contrast the value of the area for C,,E2, given by Teo et a l l 7 is much higher than would be expected from the results of Barry and El Eini15 for C&17 and C16E,,. Overall these data indicate that the comparison of minimum areas from different laboratories and involving different samples of nominally the same surfactants must be treated with considerable caution particularly for molecules with long polyoxy- ethylene chains.We concur with the conclusion of Teo et al.17 that a re-examination is desirable of the effect of oxyethylene chain length on the adsorption behaviour of carefully prepared and characterised CnE, surfactants over a wide range of values of rn. The results for low (< ca. 8) values of rn are reasonably reproducable and for C,,E, show that the increased spatial requirements of lengthening the solvated oxyethylene chain prevents close packing of n-alkyl chains perpendicular to the water/air interface and increases the average surface area occupied by each moleculeA. GELLAN AND C. H. ROCHESTER 1509 at surface saturation. The C,, part in C12E, cannot be lying parallel to the interface at high surface excess concentrations and it is unlikely that the parallel orientation prevails for C E,.The present c.m.c. and ycmc data (table 1) are consistent with the results of a more general study of the micellisation of Cl,E, (m = 2-8) surfactants by Rosen et a1.8 Increasing the hydrophilic part of the solute from E, to E, made the standard Gibbs free energy of micellisation less negative and therefore increased the c.m.c. Conversely the surface pressure zcmc was less for Cl,E, (38.0 mN m-l) than for C12E5 (41.5 mN m-l) and combined8 with the area and c.m.c. data gave standard Gibbs free energies of adsorption AGid, for Cl,E, and C,,E, at the water/air interface of - 37.9 and - 38.6 kJ mol-l, respectively. Resultant values of AGkii, - &Zds were 14.0 kJ mol-l for Cl,E5 and 15.6 kJ mol-1 for Cl,E8 (Rosen et al., obtained 12.4 and 14.8 kJ mol-l, respectively), confirming that the increment of three oxyethylene units to C12E, enhanced the ease of adsorption to form a hypothetical monolayer at minimum area per molecule and zero surface pressure relative to the ease of micellisation.Factors which might contribute to this result have been discussed before., The effects of polyethylene oxide chain length on the initial slopes of adsorption isotherms at low equilibrium concentrations have not been well investigated.l One set of isotherms for Cl,E, (m = 6, 8, 10 and 12) on Graphon showed that the extent of adsorption (in mol ern-,) was smaller the higher the value of rn for the entire concentration range studied.12 However, in accordance with the present data, Klimenko et aI.,O reported that the initial slopes of the isotherms for C,E, on acetylene black were steeper the greater the value of m.An analogy may be drawn between the isotherms in fig. 2 for C12E5 and C12E, and isotherms for a series of polyoxyethylene nonylphenols (C,4E,) adsorbed on carbon black.,' For example, the isotherm for C,#E, was steeper than that for C,#E, at low equilibrium concentrations in solution although the plateau value of the surface excess concentration was less for C&E, then for C,#E5. The average area occupied per C,,E, or C12E8 molecule at the cross-over point in fig. 2 was too small to be consistent with the surfactant molecules lying flat on the carbon surface as an estimatedI7 surface area for an adsorbed dodecyl chain alone would be 0.70 nm2. At this point, therefore, one possibility consistent with the results would be that the hydrophobic alkyl chains were orientated parallel to and interacting with the carbon surface but that the ethylene oxide segments were predominantly extending into the aqueous phase.Using a previous method of calculationz2 estimates of the areas occupied by C12E5 and C12E, lying flat on the carbon surface were 1.70 and 2.10 nm2, respectively, which would correspond to 0.70 x lop4 and 0.56 x lop4 mol g-l for C12E5 and C12E, on the present sample of Vulcan-3-G. There is a noticeable change in slope of the isotherm for C12E5 at (0.50-0.70) x lop4 mol g-' [fig. 1 (a)], for which one explanation would be that saturation of the surface by molecules adsorbed in a flat orientation was immediately superceded by the displacement of adsorbed polyethylene oxide segments from the surface by the alkyl chains of incoming adsorbate.The suggestion that not only the hydrophobic but also the hydrophilic groups in C,E, surfactants interact with carbon surfaces immersed in water was proved realistic by study of the adsorption of hexaoxyethylene glycol on G r a p h ~ n . ~ Corkill el aL3 also observed that a change in the differential heat of adsorption of C8E, on Graphon occurred at a surface coverage for which the average area occupied per adsorbed molecule approximated to the expected area if both the hydrophobic and hydrophilic parts of each molecule lay flat in contact with the surface. The present enthalpy data show that the differential1510 ADSORPTION OF SURFACTANTS heat of adsorption AET of C12E5 was constant for a range of coverages equivalent to ca.(0.2-1.7) x mol g-l [fig. 4(a)]. However, the results also show that An was more exothermic than -9.2 kJ mol-l (table 2) at low equilibrium concentrations of surfactant in solution. The two data points corresponding to < 0.2 x mol g-1 absorbed were of lower accuracy than those for higher amounts adsorbed and therefore a reliable value for AHin the very-low-coverage range could not be deduced. Despite this the results for C12E5 are essentially similar to the corresponding results for C8E63 and C8E46 on graphitised carbon, which suggests that adsorption initially involved both the hydrophobic and hydrophilic units of adsorbed molecules lying flat on the carbon surface. The present results for C12E8 would also be consistent with this possible mode of adsorption.The relative amounts of adsorption of C12E5 and C&, on carbon for equilibrium bulk concentrations above the c.m.c. are consistent with reports that increasing m in C,E,, with n constant, decreases the amount adsorbed for a particular bulk c~ncentration.~* l2 Note the similarity between the minimum average areas occupied per molecule in the carbon surface at surface saturation and the corresponding areas occupied at the water/air interface. The highest coverage recorded here for C12E5 corresponded to an average area of 0.63 nm2 per adsorbed molecule. However, surface saturation was not achieved [fig. 2(a)]. The minimum area of ca. 0.54 nm2 per molecule, estimated from data elsewhere,13 for C12E5 on carbon black compares with the minimum area of 0.56 nm2 for C12E5 at the water/air interface.The conclusion that the dodecyl groups cannot be lying flat in the surface is equally applicable for both the carbon/water and water/air interfaces when the average area per molecule is < ca. 0.70 nm2 for a C,, surfactant. The present results for Cl2E5 are therefore consistent with data for C,E, adsorbed on Sterling FTD-4, which led to the suggestion that octyl groups were partially lifting off the carbon surface at high coverages.6 A change in the differential heat of adsorption with coverage for C8E4 was ascribed to this effect.6 No similar change in AHwith coverage was observed for C12E5 [fig. 4(a)], although this is not inconsistent with the present suggestion since the highest surface coverage for which enthalpy data were measured corresponds to an area per molecule (0.68 nm2) which would allow almost all the segments of dodecyl groups in every adsorbed molecule to be flat on the carbon surface.The minimum area of 0.77 nm2 at the plateau in the isotherm for Cl,E8 would more clearly allow all the dodecyl groups to be adsorbed in a flat orientation. However, the change in value of AR for Cl,E, at a surface excess concentration of ca. 1.57 x 10-lo mol cm-2 strongly suggest that the onset of a change in the orientation of adsorbed molecules occurs at this coverage. The change, as before,6 may be ascribed to partial desorption of alkyl segments from the surface until a limiting situation is reached in which only a fraction of the alkyl segments in C1,E8 remains in contact with the carbon. The absence of this effect for C12E, with higher amounts adsorbed than those consistent with the change for C1&, leads to speculation that the increase in oxyethylene chain length from E5 to E, promotes the partial desorption of alkyl groups.A similar comparison cannot be made for C,E46 and C8EG3 because the A H data for C,E,, which exhibited no corresponding change in AE, only relate to surface coverages up to 1.9 x lo-, mol m-, whereas the second change in value of AHfor C8E4 occurred at 2.7 x lo-, mol m-,. In general the present results are in accordance with the conclusion from previous isotherm20* 2 3 9 24 and calorimetric3. 5 * data that C,E, surfactants are adsorbed on a bare graphitised carbon surface in water with both the alkyl and oxyethylene chains interacting with the carbon, but that as the surface coverage is increased oxyethylene chains are displaced from the surface by the alkyl chains of further molecules ofA.GELLAN AND C. H. ROCHESTER 1511 adsorbate. The differential enthalpy changes (Am accompanying the displacement processes for C12E5 and C,,E, are compared in table 2 with the enthalpy changes for micellisation and adsorption at the air/water interface for C12E5 and C1&8. The more endothermic values of AHmic and AH,,, for Cl2E, than for C12E5 were ascribed by Rosen et aL8 to an increase, with increasing oxyethylene chain length, in the number of hydrogen bonds between solvent water and oxyethylene chain 0 atoms, which are broken during micellisation and adsorption.The present results for adsorption on carbon are similar in that ARis less exothermic for C,,E, than for C12E5. As before6 we propose that C,E, molecules lying flat on the carbon surface at very low coverages retain some water molecules of hydration associated with the adsorbed oxyethylene segments and the hydroxyl end group in each adsorbed molecule. The enthalpy of adsorption is then dominated by exothermic contributions from the removal of alkyl chains from solution in and the interactions between the carbon surface and adsorbed alkyl chains.,'9 28 Endothermic contribution^^^ due to partial dehydration of oxyethylene groups would be at least partially compensated by the exothermic contributions resulting from the interactions between oxyethylene groups and carbon.At higher surface coverages the much less exothermic AR values (table 2) would be compatible with a model in which displacement of oxyethylene groups from the surface by increasing the excess surface concentration of alkyl groups is also accompanied by dehydration of the oxyethylene chains. The exothermic effects associated with the adsorption of alkyl groups are then largely c~mpensated~>~ by endothermic contribution^^^ due to dehydration of oxyethylene groups, both being displaced from the surface by alkyl groups and in the adsorbing molecules. The compensation was more effective for C12E8 (AR less negative) than C12E5, in agreement with the idea8 that increasing the number of oxyethylene segments in the adsorbate molecule increases the amount of dehydration which occurs.The AHvalues for C,E, (AR = 0 kJ mol-1)6 and C8E,3 are consistent with the corollary to the present proposals that reducing the alkyl chain length also reduces the exothermic contributions to AH. The proposal that C,E, surfactants are adsorbed on carbon at high surface coverages as three-dimensional associates (half-micelles) resulted from favourable comparisons between saturation adsorption values and the areas expected to be occupied by the micellar The similarity for C,,E,, between the enthalpies of micellisation and adsorption at high coverage on acetylene black has been taken as supporting this model of a d s o r p t i ~ n . ~ ? ~ ~ However, the model has been criticised on the grounds that the minimum areas occupied by the surfactants on carbon at saturation coverage are generally similar to the areas occupied at saturation in the water/air interface, where it is unlikely that hemispherical micelles exist., A further parallel between adsorption behaviour at the carbon/water and water/air interfaces is provided by the close similarity (table 2) between the present endothermic enthalpy of adsorption of C,,E, on carbon at high coverages and the corresponding enthalpy of adsorption AHWIA deduced from surface-tension data for C1&8 at the c.m.c.as a function of temperature.8 It would be more plausible to interpret the existing results for adsorption on carbon at saturation coverage in terms of a planar monolayer model, in which each alkyl chain is only partly involved in direct interaction with the carbon surface6 and the minimum area occupied by each molecule will be dependent on the stereochemical requirements of the oxyethylene chains protruding into the aqueous phase.Possible oxyethylene chain configurations are 32 and will be influenced by the degree of hydration of the chain.32333 We thank the S.E.R.C. for the award of a CASE Studentship to A.G. and Unilever Research for collaboration.1512 ADSORPTION OF SURFACTANTS I J. 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ISSN:0300-9599    

DOI:10.1039/F19858101503

出版商:RSC    

年代:1985

数据来源: RSC