摘要:

J . Chern. SOC., Faraday Trans. 1, 1982, 78,1525-1531 Kinetics of the Oxidation of Benzyl Alcohol by Tris(2,2’-bipyridine)nickel(111) Ions in Aqueous Perchlorate Media B Y DAVID F O X AND CECIL F. WELLS* Department of Chemistry, University of Birmingham, Edgbaston, P.O. Box 363, Birmingham B15 2TT Received 10th June, 1981 The oxidation of benzyl alcohol by [Ni(bipy),I3+ is first-order in both Ni”I and the alcohol. The observed independence of the second-order rate constant with acidity conforms with the following previous findings: (a) [Ni(bipy),13+ oxidizes substrates in the outer sphere without removal of a 2,2’-bipyridine from Ni”’; (b) equilibrium studies on solvent-sorting in the solvation of the proton predict that the tendency for benzyl alcohol (BzOH) to form an oxidatively inactive {H+(H,O),-, ROH) complex from ROH and (H+(H,O),} should be low.The transition-state parameters are compared for several substrates oxjdized by [Ni(bipy),13+. Following the kinetic investigation of the oxidation of various inorganic substrate ligands by [Ni(bipy)3]3+,1-3 an account is now given of the kinetic investigation of benzyl alcohol by this complex in aqueous perchlorate media. This complex participates in genuine cation + ligand redox reactions rather than in free radical oxidation, as the e.s.r. spectrum of the complex shows3 that it is a complex of the highly oxidizing Ni3+ cation and not a cation-radical complex involving NiII and a radical site on a bipyridine ligand. It was hoped that the oxidation of propan-2-01 by [Ni(bipy),13+ could be investigated kinetically to provide a comparison with the kinetic studies involving this alcohol and oxidizing aqua-~ations.~ However, although the rate of disappearance of [Ni(bipy),13+ in an excess of propan-2-01 in aqueous perchlorate media is first-order in the complex, the reaction is very slow and the observed rates at 20 O C are too close to that found for the oxidation of water by [Ni(bipy),13+ under the same conditions to provide accurate rate mea~urernents.~ At the other end of the scale, the rate of oxidation of p-benzohydroquinone by [Ni(bipy),13+ was found5 to be too fast to follow using the stopped-flow technique at 15 OC.However, the stoichiometry of this reaction is reported here, as it has a general significance for the mechanism of the oxidation of organic hydroxy compounds by this complex.Benzyl alcohol is oxidized by [Ni(bipy),I3+ at a convenient rate and, despite the restrictions imposed by the low solubility of the alcohol, a complete kinetic investigation involving the variation of [alcohol] and of [H+] was possible. EXPERIMENTAL MATERIALS [Ni(bipy),13+ was prepared as described A 5 x lo-* mol dm-, solution of [Fe(bipy),I2+ was prepared by dissolving the calculated weight of AnalaR FeSO, 7H,O and AnalaR 2,2’-bipyridine in water. Solutions of sodium perchlorate were prepared by the accurate neutralization of a solution of AnalaR HClO, with solid AnalaR sodium carbonate with boiling to expel CO, and subsequent filtration through Whatman no. 42 paper. Water was distilled 15251526 OXIDATION OF BENZYL ALCOHOL BY [Ni(bipy),13+ once in an all-glass still and AnalaR HC10, was used in reaction mixtures.Peroxy compounds were removed from the benzyl alcohol by shaking with a solution of Fe", followed by washing of the alcohol layer four times with distilled water and fractional distillation. The procedure used previously4 for the removal of carbonyl compounds from other alcohols proved to be impracticable with benzyl alcohol owing to elevation of the boiling point by the added 2,4-dini trophenylhydrazine. PROCEDURE The rates of disappearance of [Ni(bipy)J3+ were followed at 360 nm in the thermostatted cell compartment of a Unicam SP500 series 2 spectrophotometer. Water was circulated from a thermostat for the higher temperatures and a water + alcohol mixture was circulated from a cryostat for the lower temperatures.The initial [NiIIIJ was ca. 5 x lo-, mol dm-3. +p-benzoquinone reaction was measured using an excess of Ni"' and quenching with a solution containing [Fe(bipy),12+. The decrease in [Fe(bipy)32+] was determined spectrophotometrically at 522 nm and comparison with a blank solution with p-benzohydroquinone absent. The consumption ratio of the RESULTS AND DISCUSSION STOICHIOMETRY The application of the procedure for the estimation of carbonyl compounds6 used for a wide range of alcoholic materials with aqua-cations [e.g. ref. (4)] was not possible here owing to the uncertainty which exists for the extinction coefficient for benzaldehyde 2,4-dinitrophenylhydrazone in alkaline conditions and because the rate of oxidation of a dilute solution of benzyl alcohol by excess [Ni(bipy),13+ is comparable with the rate of oxidation of water by this complex.By comparison with the oxidation of simple alcohols by aqua-cations [e.g. ref. (4)], and as benzaldehyde is the principal product of the oxidation of benzyl alcohol by the permanganate a consumption ratio of IAINirrl]I/JAIBzOH]J = 2.0 would be expected. p-Benzohydroquinone also requires the removal of two electrons to produce benzoquinone, as found in its oxidation by aqua-catiom8 We found that IA[Nirrr]l/lA[QH,]1 = 2.0 0.1 (QH, = p-benzohydro- quinone) for [HClO,] = 1-5 mol dm-3, initial [NiII'] = 7.5 x mol dm-3 and initial [QH,] = 1.56 x mol dm-3; this supports the above assumption for the oxidation of benzyl alcohol in two one-electron steps by [Ni(bipy),13+. ORDERS OF REACTION Plots of log (optical density) against time for [C,H,CH,OH] + [NiIII] were always linear, showing that the reaction is first-order in [NilI1].Fig. 1 shows that the observed first-order rate contrast k , taken from the slopes of such first-order plots gives a straight line passing through the origin when plotted against the concentration of benzyl alcohol, showing that the reaction is also first-order in benzyl alcohol in 2.00 mol dm-3 HClO, at 3.00 OC. Values for k, are given in table 1, which shows that the rate is unaffected when the reaction is carried out under nitrogen. With the high yield of benzaldehyde from the oxidation by the permanganate ion7 it is unlikely that there will be attack of Nirrl on benzaldehyde when benzyl alcohol is in high excess over the Ni'II.VARIATION OF RATE WITH ACIDITY A N D TEMPERATURE Linear pseudo-first-order plots were observed for other acidities at 3.00 "C and values of k, are given in table 1 . Ionic strength was adjusted to 2.00 mol dm-3 by the addition of sodium perchlorate. Calculated values for the second-order rate constant k , given in table 1 show that k , is independent of acidity in the range 0.4- 2.0 mol dm-3 HClO, at a constant ionic strength of 2.00 rnol dmP3.D. FOX AND C. F. WELLS 1527 Linear pseudo-first-order plots were found also for 12.7, 25.3, 32.8 "C with acidity varying in the range 0.40-2.00 mol dm-, HCIO, and ionic strength maintained at 2.00 mol dm-3. Values for k , and k , are collected in table 1.The constancy of k , at constant acidity and temperature confirms that the reaction is first-order in [NiIII] and in [benzyl alcohol] and the invariance of k , at constant temperature with varying FIG. 1 5 10 15 20 1 0 [ benzyl alcohol] /mol dm-:% .-Variation of the pseudo-first-order rate constant k , with [BzOH] in 2.00 3.0 O C . mol dmP3 HCIO, at acidity confirms that the rate is independent of acidity. The mean values for k , at each temperature are also given in table 1. Fig. 2 shows that a plot of log k , against reciprocal of absolute temperature is linear and the application of the least-squares procedure gives the enthalpy of activation AH* = 77& 1 kJ mol-1 and the entropy of activation AS* = 2.0+ 1.3 J K-l mol-l. MECHANISM OF THE OXIDATION The consumption ratio found above for the oxidation of p-benzohydroquinone by [Ni(bipy),I3+, together with the evidence from the oxidation of other s~bstratesl-~ by [Ni(bipy),13+, and the consumption ratios [A[cation]l/IA[substrate]l = 2 usually found for the oxidation of alcohols by aqua-cations all suggest a two-step mechanism as in reactions (1) and (2): k2 Nil" + C,H,CH,OH -+ NilI1 + C,H,cHOH + H+ (1) with k , 9 k,.No evidence is found above for the existence of intermediate complex formation between [Ni(bipy),13+ and the alcohol and it is therefore concluded that the oxidation is an outer-sphere reaction without removal of a bipyridine from the NiIII, as found for the oxidation of H,O,,l Br-, and HN,., The absence of any kinetic effect of the hydrogen ion also supports this view, as a removal of a bipyridine from NilI11528 OXIDATION OF BENZYL ALCOHOL BY [Ni(bipy)J3+ TABLE 1 .-VALUES FOR THE PSEUDO-FIRST-ORDER RATE CONSTANT AND THE SECOND-ORDER RATE CONSTANT FOR THE OXIDATION OF BENZYL ALCOHOL BY [Ni(bipy),13+ AT IONIC STRENGTH 2.00 mol dm-l ~~ temp.[HCIO,I 1 03[BzOH] 10 k , /"C /mol dm-3 /mol dm-3 lo3 k,/ S-' /dm3 mol-' s-' 3.0 3 .O 3.0 3.0 3.0 3.0 3.0 3 .O 3 .O 3.0 3.0 3.0 3.0 12.7 12.7 12.7 12.7 12.7 12.7 12.7 12.7 12.7 12.7 12.7 12.7 12.7 12.7 25.3 25.3 25.3 25.3 25.3 25.3 25.3 25.3 25.3 25.3 25.3 25.3 25.3 32.8 32.8 32.8 32.8 32.8 32.8 32.8 32.8 32.8 32.8 2.00 15.4 0.69 2.00 30.9 1.32 2.00 62 2.76 2.00" 62 2.60 2.00 124 5.4 2.00 185 8.8 1.60 19.3 0.84 I .60 93 4.14 1 .oo 62 2.78 1 .oo 93 4.20 1 .oo 124 5.8 0.40 49.4 2.40 0.40 62 2.70 0.09 2.00 9.3 1.34 2.00 19.3 2.70 2.00 38.6 5.2 2.00 62 8.4 2.00 74 10.2 1.60 9.3 1.48 1.60 74 11.0 1 .oo 19.3 2.84 1 .oo 38.6 5.6 1 .oo 62 8.7 1 .oo 74 10.6 0.40 38.6 5.4 0.40 62 8.4 0.40 74 11.0 mean lo2 k,/dm3 mol-l s-' = 7.2k0.3 2.00 3.86 2.50 2.00 7.7 4.72 2.00 15.4 8.7 2.00 19.3 10.5 1.60 3.86 2.60 1.60 7.7 5.0 1.60 15.4 9.2 1.60 19.3 11.0 1 .oo 7.7 5.1 1 .oo 15.4 8.8 1 .oo 19.3 11.0 0.40 15.4 9.2 0.40 19.3 11.6 mean k,/dm3 mol-' s-' = 0.303 k0.020 2.00 2.35 1.51 2.00 3.86 2.49 2.00 6.2 3.80 2.00 9.3 5.7 1.60 3.86 2.53 1.60 6.2 3.90 1.60 9.3 5.7 1 .oo 6.2 4.08 1 .oo 9.3 6.0 0.40 12.3 7.7 mean k2/dm3 mol-' s-' = 0.64k0.02 mean lo2 k,/dm3 mol-l s-' = 2.24 0.224 0.214 0.223 0.210 0.218 0.238 0.218 0.223 0.224 0.226 0.232 0.243 0.218 0.72 0.70 0.68 0.68 0.69 0.80 0.74 0.74 0.72 0.7 1 0.72 0.70 0.68 0.74 3.24 3.07 2.84 2.75 3.37 3.25 2.99 2.85 3.29 2.86 2.85 2.99 3.01 6.4 6.5 6.1 6.1 6.6 6.3 6.1 6.6 6.5 6.3 a Under nitrogen.D.FOX AND C . F. WELLS 1529 3.2 3.3 3.4 3.5 3.6 3.7 lo3 KIT FIG. 2.-Variation of log (second-order rate constant) with reciprocal of absolute temperature at ionic strength = 2.00 mol dm-3. will be a~id-dependent.~'~ With its acid-independence the reaction resembles the oxidation of B r 2 and H202,1 the acid-dependence of the oxidation of HN, being ascribed to N; being the oxizable form of hydrazoic acid., The oxidation of secondary alcohols by Mn&1,47 CoIII aq, Vv aq and Ag& are all believed to be outer-sphere reactions without the involvement of intermediate cation + substrate complexes, whereas such complexes were found for the oxidation of secondary alcohols by Ceiz.4+ lo Intermediate cation + substrate complexes are believed to be involved in the oxidation of methanol by Mn:i1.l1 Following the observation12 of kinetic effects on redox reactions in aqueous solu- tions of solvent-sorting in the environment of protons, it was found that this solvent- sorting could be characterised using equilibrium measurements by a concentration quotient13 which increases markedly as the co-solvent concentration increases14 and changes the environment.It was then found4q that the relative insensitivity to changes in acidity of the rate of oxidation of secondary alcohols by Mnikl via an outer-sphere mechanism arises from a balance between the two equilibria (3) and (4) K h Mnii MnOHii + H,+, (3) Kh and K, being the concentration quotients for reactions (3) and (4), respectively, and the species (H+(H20)x-1 ROH) being unreactive towards oxidation.12 This assignment of mechanism for the outer-sphere oxidation by aqua-cations was supported by the kinetic investigation4 of the outer-sphere oxidation of propan-2-01 by Coit where Kh % Kh for Mnky: the above balance achieved with Mniy is destroyed and an acid-variation of the rate is ob~erved.~ Alternatively, this balance found with secondary alcohols and Mn;" will be destroyed producing an acid-variation in the rate if an alcohol is used which has K , @ K , for propan-2-01.As [Ni(bipy),13+ 50 FAR 11530 OXIDATION OF BENZYL ALCOHOL BY [Ni(bipy),I3t has no hydrolytic equilibrium analogous to reaction (3), the absence of an acid-variation on the rate of oxidation of benzyl alcohol is explainable if K , for benzyl alcohol is very low.Unfortunately, the solubility of benzyl alcohol in water is too low for the equilibrium methodl39 l4 involving spectrophotometric measurements on added p- nitroaniline to be used. However, the variation of K, with structure for a wide range of alcohols and carbonyl compounds shows that K, is decreased by the presence of an electron-withdrawing substituent and enhanced by the presence of an electron- releasing substituent. Although the phenyl group in phenol enhances K, by the electrons supplied to the basic oxygen atom by the n-bonding mechanism overriding the electron-withdrawing inductive effect of C6H5- as repre~entedl~ by the Taft o* function, the former will not operate at all in benzyl alcohol, allowing the latter effect to operate fully in producing this very low value for K,.TABLE 2.-cOMPARISON OF ENTHALPIES AND ENTROPIES OF ACTIVATION FOR THE OXIDATION OF SUBSTRATES BY [Ni(bipy),13+ IN AQUEOUS PERCHLORATE MEDIA A€€* AS* substrate /kJ mo1-l /J K-' mol-l H202 38+2 - 126f7 Br- 60f4 -3+11 N3 36+3 -3+10 C,H,CH,OH 77f 1 2 . 0 + 1.3 The values of AH* and AS* found for the oxidation of various sub~tratesl-~ by [Ni(bipy),13+ are compared with those for the oxidation of benzyl alcohol in table 2. The major contributions to AS* are1-, AS,*, arising from the release of oriented water following the lowering of the charge on [Ni(bipy),ln+ in the transition state and the loss of any charge on the substrate, and AS:, arising from the restriction imposed on the solvent in the transition state through any loss of a proton.Thus values of AS* for the similar reactions NiIII + Br- and NilI1 + N;, where only AS,* contributes, agree very well ; whereas AS* for NiIII + H202 has a large negative value through AS: counterbalancing AS,*. However, NilI1 + C,H5CH20H, which is very similar to NiI" + H202 (both have a radical and H+ in addition to NiII as products), has a small positive AS*. As the contributions arising from AS,* for these last two reactions should be broadly similar, the difference found suggests that there might be a different mechanism for the proton release in the two cases.It is possible that, whereas electron release and proton release are coincident with H202, they may be sequential with C6H5CH20H. In the latter case an electron may be lost first from the n-system of the phenyl group and the proton lost in a later rearrangement of charge and electron density. The restriction imposed on the solvent molecules by the diffuse charge on the phenyl group will be much less than that imposed by a proton, resulting in a higher entropy in the transition state and a smaller negative or a small positive AS* value, as observed. C. F. Wells and D. Fox, J. Chem. SOC., Dalton Trans., 1977, 1498. C. F. Wells and D. Fox, J. Chem. SOC., Dalton Trans., 1977, 1502. J. K. Brown, D. Fox, M. P. Heyward and C. F. Wells, J. Chem. SOC., Dalton Trans., 1979, 735.C. F. Wells and G. Davies, Trans. Faraday Soc., 1967, 63, 2737; C. F. Wells and M. Husain, Trans. Faraday Soc., 1970,66,679; R. Varadarajan and C. F. Wells, J. Chem. Soc., Faraday Trans. I , 1973, 69,521 ; C. F. Wells and A. F. M. Nazer, J. Chem. SOC., Faraday Trans. I , 1976,72,910; C. F. Wells and D. Fox, J. Inorg. Nucl. Chem., 1976, 38, 287.D. FOX AND C . F. WELLS 1531 D. Fox and C. F. Wells, unpublished results. ti C. F. Wells, Tetrahedron, 1966, 22, 2685. ’ K. K. Banerji and P. Nath, Bull. Chem. SOC. Jpn, 1969, 42, 2038. * C. F. Wells and L. V. Kuritsyn, J. Chem. SOC. A , 1969, 2575, 2931; 1970, 676, 1372. l o C. F. Wells and M. Husain, Trans. Furuduy Sac., 1970, 66, 2855. l 2 C. F. Wells, Discuss. Faraday SOC., 1960, 29, 219; Trans. Faraday SOC., 1961, 57, 1703, 1719. l3 C. F. Wells, Trans. Furaday SOC., 1965, 61, 2194; 1966, 62, 2815; 1967, 63, 147; Hydrogen-bonded Solvent Systems, ed. A. K. Covington and P. Jones (Taylor and Francis, London, 1968), pp. 323-324; J , Phys. Chem., 1973, 77, 1994, 1997; J . Chem. SOC., Faruduy Trans. I , 1972, 68, 993. C. F. Wells, J . Chem. SOC., Faruduy Trans. I , 1973, 69,984; 1974,70, 694; 1975,71, 1868; 1976,72, 601; 1978, 74, 636, 1569; 1981, 77, 1515; Adv. Chem. Ser., 1~79, 177, 53; G. S. Groves and C. F. Wells, unpublished results. C. F. Wells, C. Barnes and G. Davies, Trans. Faraday SOC., 1968, 64, 3069. C. F. Wells and C . Barnes, J. Chem. SOC. A , 1968, 1626; 1971, 430. l5 K . B. Wiberg, Ph-ysical Organic Chemistry (Wiley, New York, 1964), p. 415. (PAPER 1/933) 50-2

ISSN:0300-9599    

DOI:10.1039/F19827801525

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J. Chem. SOC., Faraday Trans. I , 1982, 78, 1533-1538 Kinetic Studies of the Formation of Copper(1) in Water + Acetonitrile Mixtures at 25 O C from the Reversible Reaction Cu2+ + Cuo s 2Cu-t BY D I P SINGH GILL* AND REETA SRIVASTAVA Department of Chemistry, Himachal Pradesh University, Simla 17 1005, India Received 18th June, 1981 The rate of the forward reaction of the reversible reaction Cu2+ + Cuo G 2Cu+ has been measured at 25 OC in 150 cm3 of water+ acetonitrile (H20 +AN) mixtures containing 5,10,20 and 40% (v/v) acetonitrile (AN) together with 0.032, 0.064, 0.16 and 0.32 mol dm-3 copper sulphate, 0.070 mol dmP3 H,SO, and 1, 4 and 8 g of copper crystals 2-30 pm in size. The rate of the reaction is strongly influenced by the concentrations of copper sulphate and AN and by the amount and particle size of the copper crystals.The rate of the reaction in an atmosphere of nitrogen is not affected by up to 0.47 mol dm-3 H,SO,. The reaction under investigation is reasonably fast and in most cases goes almost to completion in 3-5 h. By reversing the reaction, the backward reaction becomes faster than the forward reaction and pure copper powder starts separating out from the solution. The present investigation is thus very useful for developing a cheap and quick method for the recovery of copper from crude samples. Copper(1) salts are unstable in aqueous solutions at 25 OC.l They can, however, be stabilized in water containing nitriles, ammonia, organic bases, cyanides, halides and gases like CO and C2H4.2t Electrochemical studies by Parker and coworkers2-* have shown that electrorefining of copper by the electrolysis of copper(1) salts involves the consumption of less electricity and hence less cost as compared with the electrolysis of copper(I1) salts.The electrorefining method for the purification of copper is, however, expensive. We have developed in our laboratory a simpler, quicker and cheaper method for the purification of copper. This method involves the reversible reaction cu2+ + CUO * 2cu+. (1) Under certain conditions the forward reaction is much faster than the backward reaction and under other conditions only the backward reaction is significant. In order to discover the optimum conditions under which the forward and the backward reactions become predominant, we have measured the rates of the reactions under different conditions at 25 O C .Kinetic studies of the forward reaction are reported in this paper and those of the reverse reaction will be given in a subsequent paper. EXPERIMENTAL Doubly distilled water (distilled over acidified KMnO,) with a specific conductance of 1-3 x S cm-l was used for the preparation of all solutions. Acetonitrile (AN) (E. Merck) of 99% purity was used without further purification. Extra pure copper sulphate crystals (Sarabhai M. Chemicals) of 99.5% purity were used as received. Copper crystals stored under nitrogen gas were sieved and those particles between 2 and 30 pm in size were selected for use. For each kinetic measurement 150 cm3 of the reaction solution was used. This solution was prepared by adding AN, corresponding to 5, 10,20 and 40% (v/v), and 0.070 mol dm-3 H,SO, to the balance of distilled water.The amount of copper sulphate required to obtain 0.032,0.064, 15331534 c U 2 + + c U o $2cU+ IN WATER + ACE TONI T R I LE 0.16 and 0.32 rnol dm-3 solutions was accurately weighed and dissolved in the appropriate H,O+AN mixture. In order to remove the dissolved air from the solution, nitrogen gas (saturated with the same H,O + AN mixture) was bubbled through the solution for 20-30 min. The kinetic measurements were made at 25k0.5 OC under nitrogen gas atmosphere. The rcaction was initiated by adding 1, 4 and 8 g of copper crystals to the reaction solution and the solution was vigorously stirred. The rate of the reaction was effected drastically by changing the stirring speed and for the efficient formation of copper(1) a high stirring speed was required. Therefore, we stirred the reaction solution with a magnetic stirrer operating at a speed of 800 r.p.m.The amount of copper(1) formed as function of time was determined by KMnO, titration. The reaction kinetics in all the cases with 1 g of copper crystals was studied at 15 min time intervals while with 4 and 8 g of copper crystals 5 min time intervals were used because the reaction in these two cases was faster. In every case the first reading was, however, taken one minute after the initiation of the reaction. In all cases kinetic measurements were repeated at least twice to see the effect of heterogeneity of the copper crystals on the reaction rate. It was found that the reproducibility from two independent sets of experiments was within 10%.RESULTS AND DISCUSSION The forward reaction of the reversible reaction (1) is extremely slow in aqueous solutions at 25 OC because copper(]) salts are unstable in aqueous solutions.' This reaction in H,O +AN mixtures in the presence of a small quantity of an acid becomes very fast (the equilibrium constant changes from a value of mol dm-3 in aqueous solutions to a value of 1O1O mol dm-3 in H20+AN mixtures).2*3 This is due to the preferential solvation of Cu+ ions by AN which stabilizes copper(1) salts in H 2 0 +AN mixtures. In H20+AN mixtures the rate of the forward reaction of the reversible reaction (1) is thus significant and the rate of the backward reaction as compared with the forward reaction is negligible. We have measured the rate of the forward reaction of the reversible reaction (1) at 25 O C in 150 cm3 of H20+AN mixtures containing 5, 10, 20 and 40% (v/v) AN together with 0.032, 0.064,O.16 and 0.32 mol dm-3 copper sulphate, 0.070 mol dm-3 H2S0, and 1, 4 and 8 g of 2-30 pm in size copper crystals. The amount of copper(1) formed in mol dm-3 ( M ) as a function of time in minutes ( t ) gave parabolic plots passing through the origin at zero time in all cases, with a steep linear rise to the plots at the beginning of the reaction. The plots slowly smoothed out after 50% of the reaction time. The amount of copper(1) formed in a fixed time was a maximum in case of 40% AN and decreased in the order 40% > 20% > 10% > 5%.In cases with 0.16 and 0.32 mol dm-3 copper sulphate the reaction kinetics could not be studied in the 40% AN mixture because in this mixture the required amount of copper sulphate did not dissolve. The rate of the reaction R (= dM/dt) was determined in all cases by the empirical differentiation method, i.e. by evaluating the slope of the linear parts of the plots of M against t at the initial stage of the reaction. The values of the rate thus obtained for various systems are reported in table 1. The uncertainty in these values is 10%. A perusal of table 1 shows that the rate of the reaction is strongly influenced by the concentration of copper sulphate, AN and by the amount of copper crystals. Increasing the concentration of copper sulphate, AN and the amount of copper crystals increased the rate of the reaction. EFFECT OF PARTICLE SIZE AND THE AMOUNT OF COPPER ON THE REACTION RATE In order to study the effect of the particle size of copper on the reaction rate, kinetic measurements were also made under similar conditions using 200 mesh copper powder as well as copper crystals between 30 and l00pm in size.The rate of the reaction with 200D. S. GILL A N D R. SRIVASTAVA 1535 mesh copper powder was 5-10 times faster than the rate with 2-30 pm copper crystals and the rate of the reaction with 30-100pm copper crystals was extremely slow as compared with the rate using 2-30 pm copper crystals. Also, the rate with 8 g of copper was more than with 4 g, and with 1 g of copper it was the lowest. This is expected because in heterogeneous surface reactions the rate of the reaction depends upon the surface area of the solid phase and the larger the surface area, the faster the r e a ~ t i o n .~ ~ TABLE ACE RATE (R) OF THE FORWARD REACTION OF THE REVERSIBLE REACTION Cu2+ + Cuo e 2Cu+, IN WATER + ACETONITRILE MIXTURES AT 25 OC R/ mol dm-3 rnin-ld concentration 5% AN 10% AN 20% AN 40% AN of c u s o , /moldm-3 la 2b 3" la 2b 3" l a 2b 3" l a 2' 3' 0.032 0.50 1.37 1.89 0.55 1.67 2.43 0.71 1.72 3.00 0.73 2.57 3.86 0.064 0.58 2.96 4.58 0.79 4.66 6.20 1.85 6.09 11.70 2.72 6.33 12.50 0.16 1.91 6.67 8.29 2.82 8.00 11.60 4.17 9.47 17.60 - - - 0.32 2.41 8.57 12.90 4.10 14.00 17.00 6.67 20.00 31.30 - - - a With 1 g copper crystals; with 4 g copper crystals; with 8 g copper crystals. All these rate values have an uncertainty of & 10%.EFFECT OF THE CONCENTRATION OF H2S04 O N THE REACTION RATE Copper(1) salts are easily hydrolysed by water in neutral medium. The possibility of hydrolysis in acidic medium, however, is eliminated. Therefore, the kinetic study of the present reaction was made in the presence of an acid; dilute H2S04 was found to be the most suitable. In some cases kinetic studies were attempted without H2S0, and in all those cases precipitation occurred after a few minutes. Kinetic measurements made under similar conditions using 0.18, 0.28 and 0.47 mol dm-3 H,SO, instead of the 0.070 mol dm-3 which we used in all of our measurements showed that the rate of the reaction in an atmosphere of nitrogen gas was not much effected by the amount of H,S04 added.Dilute H,S04 has been found not to react with copper metal in the absence of oxygen.'9 In the present measurements all the dissolved oxygen was removed by bubbling oxygen-free nitrogen gas through the mixture and performing the measurements in an atmosphere of nitrogen gas; therefore, the chances of a side reaction of copper metal with H2S04 were eliminated. DEPENDENCE OF RATE O N CU2+ CONCENTRATION From table 1 it is also evident that the increased concentration of copper sulphate and hence Cu2+ increased the rate of the reaction. The dependence of rate on Cu2+ concentration has been determined by plotting log (rate) from table 1 against [Cu2+] for 1, 4 and 8 g of copper crystals in mixtures containing 5, 10 and 20% AN, [Cu2+] being the molar concentration of Cu2+.In all cases the plots were linear and the slope of these plots gave an average value of l.O+O.O5 as the dependence of the rate on [Cu2+]. This observation is confirmed in the next section. PLOTS USING DIMENSIONLESS PARAMETERS In order to find out whether the increased AN concentration changed the order of the reaction or only influenced the rate of the reaction by bringing about changes1536 c U 2 + -f- c U o $2cU+ I N WATER -k ACE TONITRI LE in the solvent composition (medium effects),s plots using dimensionless parameters in the form recommended by Powell and extended by Frost and Pearsong were made in all cases. The relative concentration, a, for the dimensionless parameter plots was calculated in each case from the ratio of the concentration of Cu2+ remaining at a particular time, i.e.Co - x, to the initial concentration of Cu2+ taken, i.e. Co. The values of x were calculated at various times from the amount of Cu+ formed using the relation Cu2+ = +Cu+. For a given amount of copper crystals (1, 4 and 8 g) with 5 , 10 and 20% AN (including 40% in cases of 0.064 and 0.032moldm-3 copper sulphate solutions) for each of the copper sulphate solutions studied, the plots using dimensionless parameters were all of the same type and comparable to the theoretical curve for the first-order rea~tion.~ This showed that the basic type of curve and hence the order of the reaction was not effected by increasing the AN concentration. Similarly, the basic type of curve was not effected when either 1, 4 or 8 g copper crystals were used.In all these cases also the plots were comparable to the theoretical curve for the first-order reaction. This showed that increasing the amount of copper crystals did not effect the order of the reaction, though an increased amount of copper crystals increased the rate of the reaction due to the increased surface area of the heterogeneous phase. For illustration the plots using dimensionless parameters for 0.16 mol dmP3 copper sulphate 1, 4 and 8 g copper crystals (2-30 pm) with 5 , 10 and 20% AN are shown in fig. 1-3. log t FIG. 1 .-Fraction of Cu2+ remaining (a) as a function of time parameter (log t) in 150 cm3 of the reaction solution consisting of 0.16 mol dmP3 CuSO, .5H,O, 1 g copper crystals (2-30 pm), x , 5 , 0, 10 and a, 20% AN and 0.070 mol drn+ H,SO, at 25 OC.EFFECT OF ACETONITRILE O N THE REACTION RATE In the light of the discussion in the previous section, it can be emphasised that for a given concentration of Cu2+ and a given amount of copper crystals an increase in the rate of the reaction with increasing AN concentration in H20+AN mixtures (table 1) can only be due to changes in the medium effects8 The medium effects influence the rate of the reaction either by change in the dielectric constant of the medium or by changes in the solvation effects.l0V1l With an increase in AN concentration the dielectric constant of the H,O + AN mixture decreases12 and one expects the formation of more ion-pairs, leading to a decrease in free Cu2+. This would lead to a decrease in the rate of the reaction.There is actually an increase in the rate with increasing AN concentration (table 1). Therefore, the increase in the rate is possibly due to the increased solvation effects.I0 Studies by Parker and coworkers2 confirm this assumption. They have reported that Cu2+ ions are preferentially hydrated and Cu+ ions are preferentially solvated by AN in H20+AN mixtures.2 Their studies have also shown that the free energy of transfer for Cu+ from pure waterD. S. G I L L A N D R. SRIVASTAVA 1537 1 I 1 0 1 2 3 0.31 log t FIG. 2.-Fraction of Cu2+ remaining (a) as a function of time parameters (log t ) in 150 cm3 of the reaction solution consisting of 0.16 mol dm-3 CuSO, .5H,O, 4 g copper crystals (2-30 pm), x , 5 . 0, 10 and 0 , 2 0 % AN and 0.070 mol dm+ H2S04 at 25 O C .0 1 2 3 log t FIG. 3.-Fraction of Cu2+ remaining (a) as a function of time parameter (log 1 ) in 150 cm3 of the reaction solution consisting of 0.16 mol dm+ CuSO, . 5H20, 8 g copper crystals (2-30 pm), x , 5 , 0, 10 and 0,20% AN and 0.070 mol dm-3 H,SO, at 25 OC.1538 c U 2 + + c U o 2cU+ I N WATER + A C ETONI TR I LE to H 2 0 + AN mixtures becomes more negative while for Cu2+ it becomes more positive as the amount of AN in the mixture increases, and the effects for Cu+ are much more pronounced than the effects for Cu2+. This indicates that increasing the concentration of AN in the mixture increases the solvation effects for Cu+ much more strongly than it decreases the hydration of Cu2+. The increased solvation effects for Cu+ stabilize these ions much more in H,O + AN mixtures with higher AN concentration and hence there is an increase in the rate of the reaction. Single-ion solvent activity coefficients for SO:-, Cu2+, Cu+ and the activated complex formed in the present reaction are not available in the literature for H20+AN mixtures and, therefore, a quantitative account of the changes of the solvation effects in terms of the solvent activity coefficientsll could not be made.But from the work reviewed by Parker1' it has been reported that due to the increased solvation effects the rate of the reaction is usually faster in dipolar aprotic solvents and in protic +dipolar aprotic solvent mixtures as compared with the rate in protic solvents. MECHANISM OF THE REACTION As the rate of the reaction was drastically increased by increasing the stirring speed, this indicated that the reaction under investigation was transport contr01led.l~ The reaction is assumed to take place not through a single step but through a number of different steps: (i) Diffusion of the preferentially hydrated Cu2+ ions from the bulk of the solution to the copper surface.(ii) Diffusion of the solvated SO:- ions from the bulk of the solution to form the negative side of the solvent-separated electrical double layer near the solid surface.14 (iii) Electron transfer from a copper atom at the surface to a Cu2+ ion to form two Cu+ cations. (iv) Solvation of Cut ions preferentially by AN to form the stable Cu+ complex. (v) Diffusion of the preferentially solvated Cu+ complex out of the electrical double layer to the bulk of the solution.Steps (iii) and (iv) are fast and should not be effected by changing the stirring speed. It is thus any one or all of the other steps which determine the rate of the reaction. We have found that under certain other conditions the backward reaction of the reversible reaction (1) becomes the significant reaction and pure copper powder starts separating out from the solution. The details of the kinetics of the backward reaction are currently under study and the results will be presented in the next paper of this series. We thank the C.S.I.R., New Delhi for a research grant to carry out this project. F. A. Cotton and G. Wilkinson, Advanced Inorganic Chemistry (Interscience, New York, 2nd edn, 1966). I. D. Macleod, D. M. Muir, A. J. Parker and P. Singh, Aust. J. Chem., 1977, 30, 1423. A. J. Parker, D. A. Clarke, R. A. Couche, G. Miller, R. A. Tilley and W. E. Waghorne, Aust. J. Chem., 1977, 30, 1661. D. M. Muir, A. J. Parker, J. H. Sharp and W. E. Waghorne, Hydrometallurgy, 1975, 1, 61, 155. K. J. Laidler, Chemical Kinetics (McGraw-Hill, New York, 2nd edn, 1965). Physical Chemistry of Process Metallurgy, Part 2, ed. G. R. St Pierre (Interscience, New York, 1961). R. G. Bates, in The Chemistry of Non-aqueous Solvents, ed. J. J. Lagowski (Academic press, New York, 1966), vol. 1 . A. A. Frost and R. G. Pearson, Kinetics and Mechanism (Wiley Eastern, New Delhi, 2nd edn, 1970), p. 14. ' R. B. Heslop and P. L. Robinson, Inorganic Chemistry (Elsevier, New York, 3rd edn, 1967). lo L. P. Hammett, Physical Organic Chemistry (McGraw-Hill, New York, 3rd edn, 1970). 'l A. J. Parker, Chem. Rev., 1969, 69, 1 . l2 A. D'Aprano and R. M. Fuoss, J. Phys. Chem., 1969, 73, 400. l 3 L. L. Bircumshaw and A. C. Riddiford, Q. Rev. Chem. SOC., 1952, 6, 157. l4 J. H. de Boer, The Mechanism of Heterogeneous Catalysis (Elsevier, Amsterdam, 1959). (PAPER 1 /995)

ISSN:0300-9599    

DOI:10.1039/F19827801533

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J. Chem. SOC., Faraday Trans. I , 1982, 78, 1539-1554 Radiation Chemistry of Xenon Trioxide, Xenate and Perxenate, and Photochemistry of Perxenate A Pulse Radiolysis and Laser Flash Photolysis Study BY ULRIK K. KLANING* Kemisk Institut, Aarhus Universitet, Langelandsgade 140, 8000 Aarhus C, Denmark AND KNUD SEHESTED Risar National Laboratory, 4000 Roskilde, Denmark AND THOMAS WOLFF Universitat-Gesamthochschule Siegen, 5900 Siegen 2 1, West Germany AND EVAN H. APPELMAN Argonne National Laboratory, 9700 South Cass Avenue, Argonne, Illinois 60439, U.S.A. Received 19th June, 1981 Unstable species containing xenon in the formal oxidation states five, Xe", and seven, Xe"", were observed by pulse radiolysis of aqueous solutions of xenon trioxide, XeO,, at pH 8-9 and of xenate, HXeO;, at pH 11-13.XevIr and species containing xenon in the formal oxidation state nine, XeIX, were observed in pulse radiolysis and flash photolysis of aqueous solutions of perxenate, HXeOi-, at pH 11-1 3. The formulae HXeO, and H,XeO;- are assumed for Xe" and XeIX, whereas the observations suggest that Xenl corresponds to three different species for which the formulae HXeO,, HXeOi- and H,XeO:- are assumed. HXeO, and H,XeO:- are formed in reactions of the hydrated electron with XeO, and HXeOi-, respectively. HXeO, and H,XeO;- are formed in reactions of the hydroxyl radical with XeO, and HXeOi- in which the hydroxyl radical adds to a ligand oxygen atom to form peroxy compounds. HXeOi- is formed in a reaction with the hydroxyl radical anion in which the hydroxyl radical anion adds to the xenon atom and by photolysis of HXeOz-: hv HXeOi- -+ HXeOi- + 0-.XeV, XeV" and XeIX and corresponding iodine species in the oxidation states four, six and eight have similar spectra and kinetics of disappearance. Estimated values of standard Gibbs energy of formation of the xenon species are used for selecting thermodynamically feasible mechanisms for one-electron reduction of perxenate and for the decomposition of perxenate in acid solution. Compounds of xenon in general show strong resemblance to the corresponding compounds of iodine.' This resemblance could extend to the properties of unstable xenon and iodine intermediates and to their role in thermal redox processes.2 Unstable intermediate species containing iodine in the formal oxidation states four, six and eight arise in the course of photolysis and radiolysis of aqueous solutions of iodate and periodate and have been described previo~sly.~-~ In the present work we investigated the radiation chemistry and photochemistry of aqueous solutions of perxenate and xenon trioxide, discussing the findings in a framework derived from the photolysis and radiolysis of iodate and periodate. 15391540 RADIATION CHEMISTRY OF XeO,, HXeO; AND HXeOi- EXPERIMENTAL NaIO,, KI, NaOH and t-butyl alcohol were Merck p.a.Perchloric acid and sulphuric acid were Merck suprapur. Ar, 0, and N,O were AGA special gasses. Crystalline sodium perxenate and an aqueous solution containing 1.1 x lo-, mol dm-3 xenon trioxide and 1 0-4-1 Ow3 mol dm-, HNO, + HClO, were prepared as previously described.6 The xenon trioxide solution was analysed iodometrically for hexavalent xenon, Xevl.Iodometric analysis of the batch of crystalline sodium perxenate indicated a formula weight of 360 & 5 corresponding to the formula Na,XeO,. 2H,O. Water used for the preparation of solutions for pulse radiolysis experiments was purified as described elsewhere.' Water used in photochemical experiments was distilled three times. All measurements were made at ambient temperature, 21 & 1 OC. Pulse radiolysis experiments were carried out using the R i s ~ HRC-linac as described.8 The dose varied between 1 and 8 krad.* The pulse duration was s. To reduce the quantity of xenon compound consumed, we used an irradiation cell that had the same optical path as that described earlier8 but that had a smaller volume (2 cm3).Steady-state photolysis, laser flash photolysis and conventional flash photolysis experiments were carried out as previously des~ribed.~~ lo Cells with I , 2, 5 and 20 cm optical paths were used in conventional flash photolysis experiments. In steady-state photolysis a cell with a 1 cm optical path was used. The laser light entering the irradiation cell was attenuated by placing brass nets in the laser beam in front of the lens which focused the light on the photolysis cell. [XeVI1'] + [Xevl] was determined by iodometric titration after converting XeVII1 to Xevl by acidification.6 The iodometric method6 was not suitable for a determination of [XeV1I1] separately in the very dilute solutions investigated here.Hence [Xevn1] was determined spectrophotomet- rically at 254 nm in solution at pH 12, taking the extinction coefficient of XevrI1 at 254 nm to be 4800 dm3 mol-1 cm-l and neglecting the contribution to the absorbance from Xevl. The pH of perxenate solutions was taken to equal 14 + log,,([NaOH] + [Na4Xe06 * 2H,0]), where [NaOH] and [Na4Xe06. 2H,O] are the concentrations of added sodium hydroxide and sodium perxenate, respectively. The pH of the xenon trioxide solutions was measured on a Radiometer PHM 52 fitted with a G.K. 2301 C electrode set. Computations were performed as previously described .5 THERMAL REACTIONS Xenon trioxide and perxenic acid and its salts are thermodynamically unstable to oxidation of water.ll However, the slowness of these reactions permits an investigation of the radiation chemistry and photochemistry of xenon trioxide and perxenate species within certain pH limits.In the present work it proved possible to study xenon trioxide in solutions with 8 < pH < 13 and perxenate in solutions with 11 < pH < 13. The rates of the thermal decompositions of alkaline xenate and perxenate solutions were found to depend strongly on the procedure employed in preparation of the solutions, indicating effects of trace impurities. Thus perxenate solutions with pH > 12 prepared according to the routine for pulse radiolysis experiments' were found to be more stable than solutions prepared under less stringent conditions. The pattern of thermal decomposition of Xevl and XeVII1 in solutions with 12 < pH < 13.3 and with [xevl] and [XeVII1] < mol dm-, was the following: Solutions containing both Xevl and XeV1I1 were less stable than solutions containing either Xevl or XeVIIJ alone. The stability of Xev' decreases with increasing pH; the stability of XevlI1 was found to pass a maximum at pH - 12.At pH 12 the stability of Xevl seems higher than that of XevrL1, since only Xevl could be detected in the solution 4 days after the preparation of the Xev1lr solution. H,O,, which may be a product in the photolysis of perxenate, reduces perxenate to elementary xenon. * 1 rad = 10+ J kg-l.U. K. KLANING, K. SEHESTED, T. WOLFF A N D E. H. APPELMAN 1541 RESULTS AND DISCUSSION The results are shown in tables 1 and 2. Table 1 lists the formulae, spectral data and values estimated for the standard Gibbs energy of formation of xenon and iodine species.Table 2 summarizes the reactions of xenon and iodine species produced by photolysis and radiolysis. PULSE RADIOLYSIS O F PERXENATE SOLUTIONS Perxenate solutions with 5 x < [Xevlll]/mol dm-3 < 5.8 x lop3 and 11 < pH < 13 were irradiated at doses varying between 1 and 8 krad. The hydrated electron, eiq, was observed in very dilute 0,-free perxenate solutions. The absorbance due to eiq decreased exponentially with time with a rate constant proportional to [XeV1*I], giving way to a transient with a spectrum depending on [XevIn]. At [Xevrr1] = 2.04 x mol dm-3, the spectrum contained a band at 560 nm and one at 370 nm. At [XeVn1] = 5 x lo-, mol dm-3, the band at 370 nm was replaced by a band at 315 nm [fig.1 (a)]. In a subsequent step, the absorbance at 13. < 370 nm decreased exponentially to zero with a rate constant that increased to a limiting value with increasing [XevlI1]. We assign the band with A, < 320 nm to a species, XeZII, formed in the reaction (1) XevrlI + e;ts + XeZ1I and we assign the band centred at 370 nm to a dinuclear species, (XeZ1IXeVIT1), in equilibrium with Xe:" and XeVIII Xe;" + Xevlrl (XeZII Xevm). (2) We assume that the primary product Xe;" has the same coordination number as the parent Xev1Ir species and that the predominant XeZ" species at 11 < pH < 13 is H3XeOi-. The rate constant of reaction (l), k,, was determined in three different ways: directly by measuring the rate of decay of the absorbance at 600 nm due to the hydrated electron in a 5.8 x lo-, mol dma3 perxenate solution, and indirectly from measurements at constant dose of the absorbances OD and ODo measured in perxenate solutions with or without addition of 0, or of BrO, (3) ODo/OD = 1 + k,[Yl/k,[XeV1ll] where [Y] stands for [O,] or [BrO,] and ki for k, or k, O,+eiq .+ 0;; k, = 1.9 x 1O1O dm3 mol-1 s-l l2 BrO,+eiq --+ BrO;+O-; k , = 2.4 x 1O1O dm3 mol-l s-l.13 (4) ( 5 ) The three values fork, thus obtained, 2.4 x lolo, 2.2 x 1O1O and 2.5 x 1 O 1 O dm3 mol-1 s-l, respectively, show satisfactory agreement.The equilibrium constant K , for equilibrium (2) was estimated graphically from a plot of OD-', the reciprocal of the absorbance measured at 320 nm, against [XeVn1]. Taking the extinction coefficient of (XeEt1XeV1I1) to be equal to zero at 13.= 320 nm, Kz may be determined from ODo/OD = l+K,[XeV1ll] (6)r-r VI R TABLE 1 .-SPECTRAL AND THERMODYNAMIC PROPERTIES OF XENON AND IODINE SPECIES absorption bands w IImax/nm ( ~ ~ ~ ~ / d m ~ mol-' cm-l) AGP/kJ mol-' 9 !2 8 < p H < 9 11 < p H < 13 8 < p H < 9 l l < p H < 1 3 . 3 5 formula formal valence -- representation 8 < pH < 9 1 1 < pH < 13 U XeV 1'' Xevl I' X e r Xe;" 1; 1: 0 z 0 275 (1000); 630 (800) - < 450 480a (800) < -120 < - 120 nil for 1 > 190 nm6 nil for II > 190 nmb 51Y 340C s 5 nil for 1. > 190 nmd nil for II > 190 nmd - 128e - 128e HXeO, 10;- 10;- 10, HXeOyb Xe0,6 10, HOOXe0, HOOIO; - 420 (- 100) nil for II > 250 nmf - < 550 < -90 305 (1700; 640 (400) - 350 (3100); 520 (400)g __ 350 (1400)g 420 - 57e CS' 350 (3000); 520 (400)g - 'V 360 (20000-2800)' 240 (5600) 3 30' nil for II > 200h - H,I 0:-g H,XeOg- 10;; H,IOiPh - 360 (- 1OOO)g 215 (5000); 280 (2800) 10,: 222 (10000)~ 3: X 366' 8 480; -41029' @ I Xelx - 580 (- 1000) 280 9 2, HOO \ (H0),/Io- - 500 (500)" - 540 (1000)" - 440 380 (1300); 550 (950) - 350 (5100)g - 520 (1100) - 510 (1200)" -410 M 350 (- 4000) a Ref.(3). Ref. (6). Ref. (11). Ref. (19). Ref. (15). f Ref. (4). g Ref. (2). Ref. (20). Ref. (21). j Ref. (22). Ref. (5).s Y O 1 x (P'Z-E) LO[ x s s D Z O I x 5'1 s soooz Z ncOI x 9'E Z sOE 1 z we01 x E'f 2 oco I x 9'E z nco I x 9'E P 60 I x 9'P s - s 8 0 1 x P'Z s L O 1 x E 21 or01 x z E 601 x f P L O 1 x 1'1 Z I 601 x 8 E'f 1-1 E'E I s.01 f ' E I E'E I I 1 < 11 < 11 < I 1 < 9 C'EI s-0 I E'E I 91-1 I - f 1-8 9 PI-1 61 901 x s 81 LOI X Z-LOI X 8'1 6- nzO1 x z-0 - 6 so1 x (S'2-z) 21 W P o 1 9'1 z so09 1 11 avo I x P'Z V I 1 nPO I x P'Z s o l X 8 > - - - I f - 11 uOI X 6'f 91 LO1 ' I or01 x E'Z L O 1 x s > 62 82 6 0 1 x f " ZE X O I - L o [ Of - E 1-21 f 1-ZI f 1-21 f 1-ZI €1-1 I E 1-1 I f 1-1 I f 1-1 I 6-8 €1-1 I f 1-1 I f 1-1 I €[-I 1 €1-1 I f 1-1 I 6-8 6-8 -1544 RADIATION CHEMISTRY OF XeO,, HXeO; AND HXeOi- E l 5 O 0 I ( a ) 1500 - - I 75 1000- - 2 m E 500- c 0 n 0 1.0 - 0.5 - I I I I I 1 250 300 400 500 6 00 700 800 X/nm FIG.1.-(a) SpectrumofeiQ adduct of XeV"' in solutions at pH 12containing[Xev1"]/mol dmP3: a, 5 x A, 2.04 x (h) I, spectrum of Xe:'; 11, spectrum of (Xe:'IXeV1'').Calculated from eqn (7) and data from fig. 1 (a). (c) Transient spectra. 0, After electron pulse irradiation of a solution at pH 12 containing mol dm-3 XeV"' and saturated with N,O at 1 atm; x , after laser pulse irradiation of an 0,-free solution at pH 12 containing 5.6 x lop4 mol dm-3 Xevllr. ODre1 are measured relative to the absorbance at 360 nm. Continuous curves are calculated from eqn (7). 0, 4.2 x mol dm-3 Xev' and 3 x where ODo is OD extrapolated to [XeVII1] = 0. We find K2 = 1600 dm3 mol-1 at pH 13. The extinction coefficients of Xe;" and (XeEII Xevnr), ~ ~ ~ ~ 1 1 and qXep X e ~ ~ l ~ ) , are related by where cVI1 is the apparent extinction coefficient and EV" = &XeV" x + &(XeF' XeV"')( 1 - x) (7) x = [Xe;I1]/([XeZII] + [(XeZII XevrIr )]> = (1 + K2[XeV111])-1.U.K. KLANING, K. SEHESTED, T. WOLFF AND E. H. APPELMAN 1545 7 1000 I 7001 ~ ~~ 400 500 600 7 1 I I I I '00 800 h/nm FIG. 2.-SpectrumofOH-adductofXeV111in solutionsat pH 12~ontaining[Xe~**~]/mol dm-3: x ,3.4 x and A, 2 x The spectra of Xe;" and (Xe;" Xevnl) were calculated from measurements of E~~~ as a function of wavelength in solutions containing [XeVm] = 4.2 x mol dm-3 and 2.04 x A transient for which Amax varies with [Xevlrl] in the range 520 < lbmax/nm < 580 remains after disappearance of Xey and (Xe2XeV1I1) (fig. 2). This transient is unaffected by addition of N,O, which quenches eCq (8) We assign the band at 580 nm to a hydroxyl adduct of XeVII1, denoted XeIx, and the mol dm-3 and are shown in fig. 1 (b).N,O + eiq + N, + 0-. band at 520 nm to a dinuclear species, denoted (XeVII1 XeIx ) XeVI1I+XeIX $(XevlllXelX). (9), (-9) (10) We assume that the addition of OH (or 0-) XeVII1 +OH (0-) --+ XeIX takes place at a ligand oxygen atom of XeVI1l to form a peroxy compound of Xevrl, the predominant species of which at 1 1 < pH < 13 is assumed to have the same charge as XezlI, i.e. (Ho)2> XeO:-. HOO The observation that the yields of XeIx and (XeVm XeIX) do not increase on addition of N,O [cf: reaction (S)] suggests that XeFI and (XezII XeVII1) react in some way to form Xelx. The following observations suggest that the mechanism for formation of XeIx consists of reactions (lo)-( 12) Xe;I1 -+ Xevl + 0- (OH) (11) (Xe;llXeV1ll) + XeV1+XeIX.(12) Reaction (1 1) is suggested by the observation that formation of 0; in 0,-containing solutions takes place in two steps: a fast step due to reaction of 0, with 0- formed as a primary species in the irradiation 0-+o, f 0, (13)1546 RADIATION CHEMISTRY OF XeO,, HXeO, AND HXeOi- and a slower step which matches the decay of absorbance assigned to XeF1 and (XeF XeVII1 1. Reaction (12) is suggested by the observation that the yields of XeIX and (XeV1l1 XeIX ) in 0,-containing solutions, in which k,,[Xevlll]/(k,,[Xevlll] + k13[02]) < lop2, are high relative to the yields of Xe;I1 and (XeFXeV1Ix). k,, and k,, were determined in the following way. Since equilibrium (2) between Xe:I1, XeVII1 and (Xey XeVII1) seems to be maintained during the decay of Xev," and (Xezl XeVII1), the apparent constant, k14, for the first-order decay of the absorbance at A < 370 nm may be expressed by k,, = ( k , , + k,, K2[XeV111])/( 1 + K2[XeV111]).(14) k,, and k,, were determined by trial and error. For K , = 1600 dm3 mol-1 we find k,, = 2.4 x lo4 s-l and k,, = 7.6 x lo4 s-l at pH 13 (fig. 3). k,, does not change measurably on changing pH from 13 to 11, which indicates that K,, k , , and k,, do not depend strongly on pH. 3 1 I 1 1 25 I 20 5 10 15 [XeVIII] / lo4 mol dm-3 FIG. 3.-First-order rate constant, kI4, of decay of transient absorbance measured at 1 < 370 nm plotted against [Xevrrl] : A, after electron irradiation and 0, after laser pulse irradiation. Continuous curve is calculated from eqn (14). The rate of formation of XelX depends on pH.The rate constant, klo, may be (15) expressed by k10 = k16y + '17(' - y ) where y = [O-]/([OH] +[0-]) = 10-pK+pH (1 + 10-pK+pH) and where k16 and k17 are the rate constants of the reactions (16) Xevrrl+ OH XeIX. (17) XevllI + 0- -+ XeIX By taking the pK of OH equal to 11 .9,14 we find from a plot of k,, against y that k,, < lo7 dm3 mol-1 s-l and k17 = 3.9 x los dm3 mol-1 s-l (fig. 4). The absorbances assigned to XeIX and (Xevrn Xexl) display complex decay kinetics.U. K. KLANING, K. SEHESTED, T. WOLFF AND E. H. APPELMAN 1547 FIG. 4.-Rate constant for formation of Xe", k,,, and [O-I/([O-l +[OHI). of Xe?, plotted against At pH < 12 the decay takes place in two steps, the slower one of which is second order. The absorbance change in the first step increases with increasing dose and decreases with increasing [XeVm].At pH > 13 the two steps merge. Furthermore, the kinetics are strongly affected by the presence of XevT. In solutions which do not contain Xevl, we suggest that XeIX and (XeVn1 XeIX) disappear via reactions ( - 9), (9), (1 8) and (19) 2XeIX -+ XeVIIT+XeV1+O, (18) 2(XeV111XeIX ) + 3XeV1I1+XeV*+0,. (19) The decay of the absorbance was studied at 550nm, at which wavelength the extinction coefficient, &IX, of both XeIX and (XeVrI1XelX) was taken equal to 1000 dm3 cm-l mol-l. Assuming that the decay at pH 12 via the slow step is due to reaction (19) alone, we find k,, = 5 x lo5 dm-3 mol-l s-l. By neglecting reactions (-9) and (19) in the decay uia the fast step, we may express the absorbance at 550 nm, OD550, as function of time t by OD550 = ~ ~ ~ l { k , [ X e ~ ~ ~ ~ ] In B/2k18 + [Xelx]t,oexp ( - k,[XevlI1]t)/B) (20) where I is the optical path of the cell and B = 1 + {2~18[Xe1x]t,o/(k,[Xev111])~ { 1 - exp (- k,[XeVIlI]t)).k , and k18 were determined from eqn (20) by trial and error to be (2.1 k0.2) x lo5 dm3 mol-l s-' respectively, [fig. 5(a)]. At pH 13, k,, k-,, k1, and k,, were determined by trial and error, using the numerically integrated rate equation corresponding to reactions (9), (- 9), (1 8) and (19), to generate corresponding values of OD550 and t . Agreement between calculated and experimental OD values was obtained with the rate constants and (1.9 k0.2) x lo7 dm3 mol-l s-l, k , = 2.5 x lo5 dm3 mol-1 s-l, k,, = 2 x 107 dm3 11101-1 s-1 k-, = 2 x lo2 s-l, k,, = 5 x lo5 dm3 mol-1 s-l and [fig.5(b) and (c)].1548 RADIATION CHEMISTRY OF XeO,, HXeO; AND HXeOi- 0.1 a 0 0.05 0 i 2 4 6 8 tl10-2 s FIG. 5.-Corresponding values of time, t , and transient absorbance, OD, at 550 nm after electron pulse irradiation of Xev1Ir solutions saturated with N,O at 1 atm. 0, pH 12; x and 0, pH 13; x , [XeV"'] = 1.67 x mol dm-3, dose 8.5 krad; 0, [XeVLrl] = 5.32 x mol dmP3, dose 1 krad; 0, [Xevrrl] = 8.33 x mol dmP3, dose 3 krad. Continuous curve (a), calculated from eqn (20). Continuous curves (b) and (c), calculated by numerical integration of rate equations corresponding to reactions (9), (-9), (18) and (19). Due to interference from the reactions H*O 20- -+ HO;+OH- it was not possible to measure &Ix directly, and it was determined in the following way.&Ix may be expressed as &Ix = where &IX' is determined directly from measurement of the corresponding values of OD,,, in N,O-saturated solutions and the dose. The quantity, a, which corrects for the fraction of 0- and OH reacting via reactions (2 1)-(23) is given by a-l = ([XelX]+[XevlllXelx I>/"O-l+ [OH]) + [OH]) / (k o[XeV"'I)> * = K l ~ ~ X e v l " ~ / ~ K ~ O ~ + ~ H ) ~ ~ O ~ ~ + + k(O-+OH)([o-] (24)u . K. KLANING, K. SEHESTED, T. WOLFF AND E. H. APPELMAN k(O-+OH) is defined as d([O-I + [OHl)/dt = k(o-+o,)([O-l+ and may be calculated from 549 k(O-+OH) = (I + 10-pK+pH)-, (2k,, 10-2pK+2pH + 2k,, 10-PK+pH + 2k,,). (26) Values for a were obtained by taking k21, k,, and k,, equal to 9 x lo8, 2.6 x 1O1O and 5.5 x lo9 dm3 mol-1 s-l, re~pective1y.l~ A value of &Ix = 1000 dm3 mo1-l.s-l was then determined graphically from a plot of a-l against &IX'.PULSE RADIOLYSIS OF XENON TRIOXIDE A N D XENATE SOLUTIONS Xevl solutions were electron pulse irradiated at varying concentration (2 x lO-*-2 x mol dm-3), pH (5-13) and dose (1-6 krad). In alkaline solution the low stability of Xevl precluded a thorough investigation, since the decomposition products 0, and XevlI1 react with 0- (OH) and eiq. Thus it proved impossible to obtain reliable results at pH > 1 1 due to the presence of XevIr1 and 0, in the solutions. Useful results were obtained with solutions saturated with N,O at 1 atm.* To correct the absorbance for the contribution from 0; formed in these solutions, the absorbance measured 150 ps after the pulse was substrated from that measured 10 ps after the pulse.The corrected absorbance is assigned to a species XeFI, formed in the reaction Xevl + 0- (OH) -+ XeV,I1. ( - l l a ) The spectrum obtained upon irradiation of a freshly prepared alkaline, N,O-saturated Xevl solution resembled that obtained with a dilute Xevlll solution that did not contain N,O [fig. 1 (c)]. The disappearance of Xe:" in a freshly prepared N,O-saturated Xevl solution proceeded at a much lower rate than in solutions that also contained XevlI1. The rate of disappearance of XeV,I1 increases on addition of XeVIII and the kinetics change to a first-order process with a rate constant similar to that measured for the decay of XeV," (fig. 3).We assume that the following reactions take place XezI $ XeV1+ 0- (OH) ( l l a ) ( - l l a ) XeF1+XeV1ll+(Xe,V1lXeV1ll) (2 a) (XeilI XevIr1) -+ XeV1+XeIX. ( 1 2 4 k-lla was determined from measurements of the pseudo-first-order rate constants for the growth of absorbance at 320 or 600 nm in solutions saturated with 1 atm N,O and containing Xevl in varying concentrations. kllU increases with decreasing pH. We take k-lla to be given by k - l l U = k28y+k29(1 - y ) (27) (28) (29) where y = [O-]/([O-]+[OH+]) and where k28 and k,, are the rate constants of Xevl + 0- -+ XezI Xevl + OH -+ Xe,VI1. k,, and k,, were determined graphically from a plot of k-,,, against y (fig. 4). We find k,, = 1.3 x lo9 dm3 mol-l s-l and k,, < 5 x lo7 dm3 mol-l s-l.At pH > 1 1 the * 1 atm = 101 325 Pa.1550 RADIATION CHEMISTRY OF XeO,, HXeO, AND HXeOi- predominant Xevl species is HXeOl6. Accordingly, we assume that XeV,I1 is formed as the pentacoordinated species HXeO:-. Upon irradiation of 0,-free XeO, solutions at pH 8-9, a transient absorbance with maxima at 275 and 625 nm is observed (fig. 6). In solutions saturated with 1 atm N,O, in which eiq is scavenged by N,O, a very weak absorbance is observed to grow in at 400-450 nm with a rate constant equal to 107-108 dm3 mol-l s-l. We assign the I I 1 I 1 I I 300 400 500 600 700 8 00 h/nm mol dmP3 Xe"' solution at pH 8.6. FIG. 6.-Spectrum of transient absorbance, OD, after electron pulse irradiation (dose 2 krad) of an 0,-free transient absorbance observed in N,O-containing solution to an OH-adduct of XeO,, Xeyl, which may be visualized as a peroxy compound of XeV, HOOXeO,, and we assign the absorbance in N,O-free solution mainly to the XeV species, HXeO,, which is formed in the reaction Xevl + eCq -+ XeV.(30) The extinction coefficients E~ of the species XeV at the two maxima J. = 275 and 625 nm are 1050 and 800 dm3 mol-1 cm-l, respectively. The transient absorbance assigned to the XeV species decays rapidly in an apparent second-order process with a rate constant equal to 8 x lo9 dm3 mol-1 s-l. Addition of t-butyl alcohol, which reacts rapidly with OH, did not change the transient spectrum. The rate of decay, however, became smaller and the kinetics changed to a first-order process. We suggest, therefore, that the observed second-order decay of transient absorbance is also due to reactions of the XeV species with OH and with XeV,I1.PHOTOLYSIS OF PERXENATE SOLUTIONS The steady-state quantum yield for disappearance of Xevrrl at pH 12 and 13 in 2 x lop4 mol dm-, 0,-free Xevlrl solutions was 0.1. However, due to lack of repro- ducibility we were unable to determine the stoichiometry of the overall photochemical process. Laser flash photolysis of 0,-free solutions containing 2 x < [Xevlll]/mol dm-, < lo-, and 0 < [NaOH]/mol dm-, < 0.1.U. K. KLANING, K . SEHESTED, T. WOLFF A N D E. H. APPELMAN 1551 resulted in a transient increase of absorbance in the wavelength region investigated (360 < A/nm < 600).The growth of the transient absorbance matched the shape of the laser pulse and was proportional to the pulse intensity.The spectrum and the kinetics of disappearance of the transient absorbance resemble the spectra [fig. 1 (c)] and kinetics of the disappearance of Xe;" and Xe;". In solutions saturated with 0, at 1 atm pressure, 0; was observed to form in two steps. These steps were kinetically first order. The first fast step is attributed to reaction (13), where 0- and Xe;I1 are formed in the primary process hv Xevrll -+ XeLII+ 0-. (31) The second step matches the disappearance of Xe;" in 0,-free solution and is attributed to reaction (1 1 a). The present flash photolysis apparatusg did not permit observation of Xe;I1. A transient increase of absorbance due to XeIX and (XevIn, XeIX) was observed in the region 400-800 nm.0; formed in 0,-containing solutions at pH 12 disappeared with concomitant formation of XeIX and (XeV1l1, XeIx). The rate of decay of 0; matched the rate of formation of XeIX and (XeVII1, XeIX). We suggest that the decay of 0; proceeds via reaction (- 13) as 0- (OH) is depleted by reaction (10). A decrease of absorbance taking place in two steps was observed at A -= 330 nm both in 0,-free and in 0,-containing Xevrll solutions at pH 12. No change in spectrum was detected and the absorbance changes were therefore attributed to a decrease in [XevrrI]. An explanation of the slow step was not possible. However, the following findings suggest that the fast step, in which the decrease of absorbance matched the profile of the flash, may be assigned to reaction (3 1) as the only primary photochemical process.We found that the yield of 0; in an 0,-saturated 2 x mol dm-, XeVII1 solution was 1.7 times the initial decrease of XeVII1 content, suggesting that the source of 0; is 0- formed by the primary reaction (3 1) and by the secondary reaction (1 1 a), which competes with reaction (12a) through the equilibrium (2a) (equilibrium constant = 1600 dm3 mol-l). SUMMARY O F REACTIONS A N D GIBBS ENERGIES OF FORMATION OF IODINE A N D XENON SPECIES Table 2 shows that XeV, Xevrl and XeIX species react very similarly to IIV, Ivl and IV1I1 species. Note, however, that the intervalence complexes (Xevl XeVII1), (XeVII XeVII1 1 and (XeVII1 XeIx) are more stable than the corresponding iodine species. The reactivities of 0- and OH are notably different.0- reacts rapidly with HXeO; and 10; and slowly with HXeOi- and H,I0,3-. OH reacts rapidly with HXeOi- and H,IOi- (HJOi-) and slowly with XeO,, HXeO; and 10;. We suggest that 0- and OH add to the Xe atom of HXeO; and that 0- adds to the I atom of 10; to form, respectively, the XeilI and I:] species HXeOi- and 10;- (table 1). Furthermore, we suggest that OH adds to a ligand oxygen atom of XeO, and 10; to form the Xei1I and 1; species, which may be visualized as the peroxy compounds HOOXe0, and HOOIO;, which contain xenon and iodine in the oxidation states five and four, respectively, XeO, +OH -+ HOOXe0,. Similarly, we suggest that OH reacts with HXeOi-, H,IOi- and H,IO;- with formation of the XeIX and IV1I1 species which contain xenon and iodine in the oxidation states seven and six, respectively.1552 It has previously been suggested that 10; reacts with OH to form 10, via the electron-transfer reaction 10; +OH -+ 10, + OH-.,$ However, an estimate of the standard Gibbs energy of formation of 10,, AGP(IO,) = 190 kJ mo1-1,2 indicates that this process is not thermodynamically feasible.We suggest that the reactions of eLq with XeO, and 10; are addition reactions. Thus we suggest that the species 10:- and HIO; result from the reaction of eLq with 10;. It has previously been suggested that the species 10, is formed in the reaction of eGq with IO;., The present measurements do not reveal any difference between Xe:" and Xe;". The cause may be that Xe:" and Xe;" interconvert rapidly or that XeV" species, despite differences in structure, have indistinguishable spectra and values of the rate constants (1 1) for the reactions H,XeOi- -+ H,XeO; + 0- (1 1 4 That k,, is similar to k,,, may be rationalized by assuming that the changes in Gibbs (33) energy of the reactions (34) are similar.Such similarities of spectra and rate constants exist for the analogous Ivl species. RADIATION CHEMISTRY OF XeO,, HXeO, AND HXe0,3- HXeOg- -+ HXeO, + 0-. HXeOg- + H 2 0 -+ H,XeOi- HXeO; + H 2 0 -+ H,XeO;. The values of AGP listed in table 1 are estimated as follows. By taking AGP(I0;) = - 128 kJ m ~ l - ~ , ' ~ AGY(Xe0,) = 515 kJ mol-l,ll AGf(0H) = 36 kJ mol-l l6 and AGY < 0 for reaction (32) and the corresponding reaction of OH with IO;, we find upper limits for and for AGp(HOOXe0,) = 550 kJ mol-l AGP(HOOI0;) = -90 kJ mol-l.Similarly, by taking AGf(HOOXe0,) - AGf(HXe0,) = AGP(HOOI0;) - AGP(HI0;) = AGP(H,O,),, - AGP(H,O), = 103 kJ mo1-1,16 we find upper limits for AGP(HI0;) = - 190 kJ mo1-I and for AGP(HXe0,) = 450 kJ mol-l. By using Pauling's rule for estimation of acidity constants17 we find the upper limit for AGP(IO:-) = - 120 kJ mol-l. AGP(HXeOg-) may be estimated as previously described for AGP of the corres- ponding Ivl species. By taking k - , , , / k l l , z 5 x lo4 mol dm-, equal to the equilibrium constant of the equilibrium (35) HXeO, + 0- $ HXeOg- and by taking AGP(HXe0;) = 340 kJ mol-l l1 and AGP(O-) = 105 kJ mo1-',16 we find AGP(HXe0,2-) = 420 kJ mol-l. Pauling's rule17 for estimation of acidity constants was used to estimate changes in AGP on protonation of Ivl species.2 However, since it seems that Pauling's rule may not be valid for xenon specie^,^ we make the following estimates.From the fact that k,,, does not depend on pH for 1 1 < pH < 13, we estimate that pK of HXe0:- is > 13. Furthermore, by taking the pK of H,Xe05 andu. K. KLANING, K. SEHESTED, T. WOLFF AND E. H. APPELMAN 1553 H,XeO; to be equal, respectively, to the pK of H,XeO, and H,XeOi-, we find for AGP of H,XeO, and H,XeO; the values 320 and 360 kJ mol-l, respectively. By taking AGP = 0 for hydration of HXe0;- [reaction (33)] and AGP(H,O), = -237 kJ m ~ l - ~ , ~ ~ we find AGP(H,XeO;-) z 180 kJ mol-l. Further, by taking the pK of H,XeO, and H,XeO; to be equal to 6 and 10, respectively, we find for AGF of H,XeO, and H,XeO; the values 90 and 130 kJ mol-l, respectively.IV1I1 and XeIX are assumed to be the peroxy compounds H,XeO;-, H,IO;- and H,IO; of six-coordinated Ivl and XeVI1. By taking AGP(XeIX)- AG,O(XeVII) = AG F(Ivlll)- AGP(IV1) = AGP(H,O,),, -AGF(H,O)l, we find the values for AGP(XeIx) and AGP(IV1I1) shown in table 1 . Note that the AGP values of xenon species show the same pattern as do those of iodine species. This is in accordance with the observation that the reactivities of xenon and iodine species are similar (table 2). Thus the values given in table 1 imply that the mechanism for a one-equivalent reduction is the same for IV1I and XeVII1. The suggested mechanism consists of a transfer of an electron to XevlI1 (Iv"), followed by either decomposition of XeVI1 (Iv1) to Xevl (Iv) and OH (0-) or transfer of an OH radical from XeVI1 (Iv1) to the reductant., Perxenate decomposes rapidly to xenon trioxide in acid solution.The reactions H,XeO, -+ H,XeO, + OH H4Xe06 + H20 3 H,XeO, 4- OH (36) (37) have been suggested as the primary step.18 However, according to our present estimates of AG: of XeVI1 species, reactions (36) and (37) are not thermodynamically feasible. Instead, we tentatively suggest the reaction 2H,XeO, -+ H,XeO,+ H,XeO, (38) for which AGO may be negative (table 1) and by which OH is transferred from one XeVm to another XeVII1 to form a five-coordinated XeVII species and the six-coordinated peroxy XeVII species H,XeO,. In analogy to the reaction of Ivl with the next step may be a transfer of OH H,XeO, + H,XeO, -+ H,O + XeO, + H,XeO,.(39) The final step may then be a decomposition of H,XeO, [reactions (18) and (19)]. We thank Arne E. Nielsen, Medicinsk-Kemisk Institut, University of Copenhagen, for pointing out that the bubble formation occasionally observed in our laser flash photolysis experiments was probably due to the presence of particulate material causing heterogeneous gas nucleation. Financial support from Statens Naturvidenskabelige Forskningsriid to one of us (U. K. K.) is gratefully acknowledged. N. Bartlett and F. 0. Sladky, in Comprehensive Inorganic Chemistry (Pergamon Press, Oxford, 1973), vol. I , p. 213. U. K. Klaning, K. Sehested and T. Wolff, J . Chem. Soc., Faraday Trans. 1, 1981, 77, 1707. F. Barat, L. Gilles, B. Hichel and B. Lesigne, Chem. Commun., 1971, 187; J. Phys. Chem., 1972, 76, 302. Y. Tendler and M. Faraggi, J . Chem. Phys., 1973, 58, 848.1554 RADIATION CHEMISTRY OF XeO,, HXeO, AND HXeOi- U. K. Klaning and K. Sehested, J. Chem. Soc., Faraday Trans. I , 1978, 74, 2819. E. H. Appelman and J. G. Malm, J . Am. Chem. Soc., 1964, 86, 2141. E. Bjergbakke, in Manual of Dosimetry, ed. N . W. Holm and R. J. Berry (Marcel Dekker, New York, 1970). H. C. Christensen, G. Nielson, P. Pagsberg and S . 0. Nielsen, Rev. Sci. Znstrum., 1969, 40, 786. U. K. Klaning, J. Chem. Soc., Faraday Trans. I , 1977, 73, 434. lo T. Wolff, J. Phorochem., 1979, 11, 215. l1 J. G. Malm and E. H. Appelman, Atomic Energy Review (International Atomic Energy Agency, Vienna, 1969), vol. VII, no. 3. l 2 Selected Specific Rates of Reactions of Transients from Water in Aqueous Solution: Hydrated Electron (US. Department of Commerce, National Bureau of Standards, Washington, D.C., 1973 and 1975). l 3 K. J. Olsen, K. Sehested and E. H. Appelman, Chem. Phys. Lett., 1973,19,213. E. H . Appelman, S. Gordon, U. K. Klaning and W. Mulac, unpublished results. l4 Reactivity of the Hydroxyl Radical in Aqueous Solution, NSRDS-NBS46 (U.S. Department of Commerce, National Bureau of Standards, Washington, D.C., 1973). l5 Selected Values of Chemical Thermodynamic Properties, Technical Note 270-3 (US. Department of Commerce, National Bureau of Standards, Washington, D.C., 1968). l6 G. Milazzo and S. Caroli, Tables of Standard Electrode Potentials (John Wiley, New York, 1978). l7 L. Pauling, General Chemistry (W. H. Freeman, San Francisco, 3rd edn, 1970). Is E. H. Appelman and M. Anbar, Znorg. Chem., 1965, 4, 1066. l9 A. Treinin and M. Yaacobi, J . Phys. Chem., 1964, 68, 2487. 2o C. E. Crouthamel, H. V. Mech, D. S. Martin and C. V. Banks, J . Am. Chem. Soc., 1949, 71, 3031. 21 C. E. Crouthamel, A. M. Hayes and D. S . Martin, J. Am. Chem. Soc., 1951, 73, 82. 22 G. K. Johnson, P. N. Smith, E. H. Appelman and W . N. Hubbard, Znorg. Chem., 1970, 9, 119. (PAPER 1/1001)

ISSN:0300-9599    

DOI:10.1039/F19827801539

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J. Chem. Soc., Furaday Trans. I, 1982, 78, 1555-1560 Structural Analysis of Molten LiBr BY HIDEO OHNO* AND KAZUO FURUKAWA Molten Materials Laboratory, Division of Nuclear Fuel Research, Japan Atomic Energy Research Institute, Tokai-mura, Ibaraki 3 19- 1 1, Japan AND KAZUO IGARASHI AND JUNICHI MOCHINAGA Department of Synthetic Chemistry, Faculty of Engineering, Chiba University, Chiba-shi 260, Japan Received 3rd July, 1981 The structure of molten LiBr has been studied by X-ray diffraction analysis and compared with that generated by computer simulation. The position of the first peak in the correlation function g(r) was found to be 0.265 nm by experiment; this value is slightly higher than that obtained by computer simulation (0.24 nm). The coordination number of the first nearest-neighbour was 3.8, which is in good agreement with the value estimated from Furukawa's relationship.Levy et a1.l have studied the structure of molten LiBr by the X-ray diffraction method and reported that the coordination number of the nearest Li-Br pair was 5.2. On the other hand, Furukawa2 has reported that the first coordination number of molten alkali-metal halides, n,, can be evaluated from a proportional relationship using the experimental values of other quantities. His estimated value of n, was ca. 4. Since the n, value gives us information about the molten structure comparable with that obtained from the position of the first peak, r,, of the radial distribution function (r.d.f.), precise values are desirable. In the present study, the r.d.f.of molten LiBr was determined by X-ray diffraction and the results are compared with those reported previous1y.l EXPERIMENTAL Experimental procedures and analysis of observed intensities are identical to those described in our previous paper.3 X-ray diffraction measurements were performed on a 8-8 X-ray diffractometer having a curved graphite crystal in the path of the diffracted beam. Mo Ka (A = 0.071 07 nm) radiation diffracted at the free surface of the melt was then monochromatized. The observable range of scattering angle (0) was 3 < < 35, corresponding to the range 9.3 < Q/nm-l < 101.4 (Q = 4 nssinO/A). The slit systems used were go-+' and lo-lo for low (3 < S/O < 15) and high (12 < S/" < 35) scattering angles, respectively. The minimum total count per datum point at intervals of 0.25O was accumulated to lo4 at low scattering angles and to 2 x lo4 at high scattering angles.The temperature of the melt was controlled at 570 f 5 "C throughout the measurements. The observed intensities were corrected for polarization and absorption in the melt in the usual manner. The background was subtracted from the measured intensities, so that the difference between the scaling factors derived both by the high-angle region method and by 15551556 STRUCTURAL ANALYSIS OF MOLTEN LiBr the method of Krogh-Moe and Norman was within 0.02%. The r.d.f., D(r), the correlation function, g(r), and the reduced intensity Q * i(Q) are given by the following expressions: g(r) = 1 +X(Km)2/(27t2gor) JoQmax Q.i(Q)sin(rQ)dQ (1) D(r) = 4nr2g0g(r) (2) i(Q> = s(Q)- 1 (3) m m and where po is the number of stoichiometric units per unit volume (lo3 nmP3), Km the effective electron number in the atom m,f,(Q) the independent atomic scattering intensity and I,C,Oh(Q) the total coherent intensity.The parameters used in the analysis are given in table 1. The sample (of analytical reagent grade) was dehydrated by heating at 500 OC under continuous evacuation for ca. 24 h and then purified argon gas was bubbled into the melt for a few hours. TABLE PA PARAMETERS USED IN THE CALCULATION OF THE R.D.F. OF MOLTEN LiBr ~~ ~ ~~ temperaturePC 570 density/ lo3 kg m-3 2.516 effective electron number: K,i 2.168 KBr 35.832 Po 0.017 47 Qmaxlnm-’ 100 RESULTS A N D DISCUSSION Fig. 1 shows the observed reduced intensity curve Q - i(Q) of molten LiBr at 570 OC.The r.d.f., D(r), the function D(r)/r, and correlation function, g(r), are shown in fig. 2. The experimental numerical structure values, Q - i(Q), S(Q) and g(r) are given in tables 2 and 3, respectively. Levy et a1.l have reported an rl value in D(r) of 0.268 nm. The rl value depends on the choice of D(r), D(r)/r and g(r), and the rl values from these three curves obtained in this work are 0.269, 0.267 and 0.265 nm, respectively, as shown in fig. 3. The rl value in D(r) obtained is in good agreement with that reported by Levy et a2.l Computer simulation by Monte Carlo (MC) or molecular dynamics (MD) has been applied to many molten alkali-metal halide+ and the structural properties reported are in good agreement with those derived from experiments by X-ray and neutron diffraction.We have previously concluded that rl values of the r.d.f. of molten alkali-metal halides calculated by computer simulation are always shorter than those found from experiments6 Lantelme et al.’ have carried out a computer simulation by molecular dynamics for molten LiBr and gave an rl value of 0.24 nm for g(r), which is slightly lower than that for g(r) obtained by X-ray diffraction (0.265 nm). The difference (0.025 nm) exceeds the experimental uncertainty, which is < 0.0 1 nm. As discussed previously,6 this small difference seems to be due to the non-spherical deformation of the electron shell in the area where ions are in contact with each other, and the polarization of the ion in the model employed in computer simulations has not sufficiently been takenH.OHNO, K . FURUKAWA, K. IGARASHI A N D J. MOCHINAGA 1557 i . 5 i c 0.5 h 01 . o 01 v .- - 0.5 - { 0 I I I I I I I 1 I I I 1 I I I I 2 4 6 0 10 Q/ 10 nm-' FIG. 1.-Reduced intensity curve of molten LiBr at 570 O C . 3 , "0 , Oi2 , 0,.4 , 0,6 I 0,8 , 1 .[ r/nm FIG. 2.-The functions D(r), D(r)/r and g(r) for molten LiBr at 570 O C .1558 STRUCTURAL ANALYSIS OF MOLTEN LiBr 0 25 0 28 r/nm FIG. 3.-First-peak positions of D(r), D(r)/r and g(r) for molten LiBr: (a) D(r), 0.269 nm; (b) D(r)/r, 0.267 nm; (c) g(r), 0.265 nm. TABLE 2.-NUMERICAL VALUES OF THE STRUCTURE Q . i(Q) AND S(Q) OF MOLTEN LiBr AT 570 O C a 1.0 -0.841 1.1 -0.904 1.2 -0.962 1.3 -0.985 1.4 -0.984 1.5 -0.721 1.6 -0.071 1.7 0.728 1.8 1.354 1.9 1.558 2.0 1.289 2.1 0.709 2.2 0.083 2.3 -0.367 2.4 -0.571 2.5 -0.594 2.6 -0.555 2.7 -0.524 2.8 -0.493 2.9 -0.406 3.0 -0.225 3.1 0.031 3.2 0.294 3.3 0.489 3.4 0.576 3.5 0.556 3.6 0.455 3.7 0.302 3.8 0.121 3.9 -0.066 4.0 -0.230 0.159 0.178 0.198 0.242 0.297 0.519 0.956 1.428 1.753 1.820 1.645 1.338 1.038 0.840 0.762 0.762 0.787 0.806 0.824 0.860 0.925 1.010 1.099 1.148 1.169 1.159 1.126 1.082 1.032 0.983 0.943 4.1 4.2 4.3 4.4 4.5 4.6 4.7 4.8 4.9 5.0 5.1 5.2 5.3 5.4 5.5 5.6 5.7 5.8 5.9 6.0 6.1 6.2 6.3 6.4 6.5 6.6 6.7 6.8 6.9 7.0 7.1 -0.344 -0.391 -0.378 - 0.328 - 0.264 -0.196 -0.1 18 - 0.022 0.088 0.192 0.264 0.289 0.270 0.221 0.161 0.099 0.036 - 0.026 - 0.079 -0.113 -0.122 -0.1 13 - 0.098 - 0.089 - 0.088 - 0.088 - 0.077 - 0.052 -0.016 0.019 0.045 0.916 7.2 0.907 7.3 0.912 7.4 0.925 7.5 0.941 7.6 0.957 7.7 0.975 7.8 0.995 7.9 1.018 8.0 1.038 8.1 1.052 8.2 1.055 8.3 1.051 8.4 1.041 8.5 1.029 8.6 1.018 8.7 1.006 8.8 0.996 8.9 0.987 9.0 0.981 9.1 0.980 9.2 0.982 9.3 0.984 9.4 0.986 9.5 0.986 9.6 0.987 9.7 0.989 9.8 0.992 9.9 0.998 10.0 1.003 1.006 0.056 0.058 0.055 0.053 0.05 1 0.051 0.049 0.045 0.035 0.019 - 0.002 - 0.02 1 - 0.033 - 0.036 - 0.033 - 0.029 - 0.03 1 - 0.038 - 0.042 - 0.036 -0.019 0.003 0.021 0.030 0.027 0.020 0.013 0.009 0.006 1.008 1.008 1.007 1.007 1.007 1.007 1.006 1.006 1.004 1.002 0.999 0.997 0.996 0.996 0.996 0.997 0.996 0.996 0.995 0.996 0.998 1 .ooo 1.002 1.003 1.003 1.002 1.001 1.001 1.001 a Q is in units of lo-' nm-l.H.OHNO, K . FURUKAWA, K. IGARASHI A N D J. MOCHINAGA 1559 into account.The effect would be relatively large in a molten salt composed of ions of different charge densities such as LiBr and LiI. The n, value gives us comparable information about the molten structure as does the r1 value. The following three methods have usually been used to determine n, values3 (a) symmetrical rg(r), (b) symmetrical rzg(r) and (c) integration to the first minimum in D(r). Usually, although not always, these methods will result in progressively higher numerical values, i.e. (nl)a < (n,), < (nl)c. The experimental n, values of molten LiBr obtained in this work for methods (a) and (c) are 3.8 and 4.1, respectively; these are smaller than that reported by Levy et a1.l (5.2). However, the n, value in this experiment is close to the results from computer simulation, 4.27.' The reason for the discrepancy between this work and the results of Levy et al.is due to the non-negligible amount of ghost below r = 0.2 nm present in the latter case; this will affect the n, value. The first coordination number can also be evaluated from the following proportional expression2 n:alc = n,( Vz/ V 3 ( r : / r 3 3 where Vz and V$ are the molar volumes of the solid and liquid at the melting point, rs and ri are the nearest-neighbour distances of the Li-Br pair in the solid and liquid, respectively, and n, is the coordination number of the nearest Li-Br pair in the solid. Using the following values, V$ = 28.03 x m3 mol-l, rs = 0.285 nm, r: = 0.265 nm and n, = 6, one obtains nfalc = 3.9. Thus, as in all (6) m3 mol-l, Vd = 34.3 1 x TABLE 3.-NUMERICAL VALUES OF THE STRUCTURE g(r) OF MOLTEN LiBr AT 570 *Ca 0.21 -0.012 0.22 0.027 0.23 0.130 0.24 0.283 0.25 0.440 0.26 0.541 0.27 0.544 0.28 0.450 0.29 0.307 0.30 0.189 0.31 0.164 0.32 0.264 0.33 0.477 0.34 0.763 0.35 1.069 0.36 1.356 0.37 1.599 0.38 1.783 0.39 1.897 0.40 1.935 0.41 1.896 0.42 1.792 0.43 1.647 0.44 1.487 0.45 1.336 0.46 1.205 0.47 1.094 0.48 0.49 0.50 0.51 0.52 0.53 0.54 0.55 0.56 0.57 0.58 0.59 0.60 0.61 0.62 0.63 0.64 0.65 0.66 0.67 0.68 0.69 0.70 0.71 0.72 0.73 0.74 1 .ooo 0.920 0.855 0.808 0.776 0.757 0.745 0.736 0.733 0.738 0.753 0.778 0.807 0.835 0.859 0.881 0.902 0.925 0.950 0.977 1.005 1.035 1.066 1.099 1.128 1.150 1.160 0.75 0.76 0.77 0.78 0.79 0.80 0.8 1 0.82 0.83 0.84 0.85 0.86 0.87 0.88 0.89 0.90 0.9 1 0.92 0.93 0.94 0.95 0.96 0.97 0.98 0.99 1 .oo 1.157 1.146 1.131 1.116 1.101 1.085 1.064 1.039 1.012 0.988 0.969 0.956 0.948 0.943 0.938 0.934 0.93 1 0.93 1 0.935 0.941 0.949 0.959 0.970 0.982 0.994 1.005 a r is in units of nm.1560 STRUCTURAL ANALYSIS OF MOLTEN LiBr molten alkali-metal halides except LiI,2 the n, value of molten LiBr estimated from eqn (6) is in reasonable agreement with the experimental value obtained by X-ray diffraction analysis.CONCLUSIONS (I) The position of the first peak in correlation function g(r) was obtained as 0.265 nm, a slightly higher value than that obtained by computer simulation (0.24 nm). (2) The coordination number of the first nearest-neighbour was found to be 3.8, in good agreement with that calculated from Furukawa's relationship. H. A. Levy, P. A. Agron, M. A. Bredig and M. D. Danford, Ann. N . Y. Acad. Sci., 1960,79, 762. K. Furukawa, Discuss. Faraday Soc., 1962, 32, 51. H. Ohno and K. Furukawa, J . Chem. SOC., Faraday Trans. I , 1978, 74, 795. M. J. L. Sangster and M. Dixon, Adv. Phys., 1976, 25, 247. L. V. Woodcock, Advance in Molten Salt Chemistry, ed. J. Braunstein, G. Mamantov and G. P. Smith (Plenum Press, New York, 1975). K. Furukawa and H. Ohno, 3rd Int. Symposium on Molten Salts (The Electrochemical Society, Princeton, N.J., 1980), extended abstracts, vol. 80-2, p. 1587; Proc. 3rd Int. Symposium on Molten Salts (The Electrochemical Society, Princeton, N.J., 1981), vol. 81-9, p. 36. ' F. Lantelme, P. Turq and P. Sochofield, J. Chem. Phys., 1977, 67, 3869; F. Lantelme and P. Turq, Mol. Phys., 1979, 38, 1003. C. J. Pings, Physics of Simple Liquids, ed. H. N. V. Temperley, J. S. Rowlinson and G. S. Rushbrooke (North Holland, Amsterdam, 1968), chap. 10. (PAPER 1/1058)

ISSN:0300-9599    

DOI:10.1039/F19827801555

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J. Chem. Soc., Faraday Trans. I, 1982, 78, 1561-1567 Tube Electrode and Electron Spin Transient Signals B Y w. JOHN ALBERY* AND RICHARD G. Resonance COM P T O N ~ Department of Chemistry, Imperial College of Science and Technology, South Kensington, London SW7 2AY Received 6th July, 1981 A tube electrode is positioned so that the electrode is just above an electron spin resonance cavity; radicals generated on the electrode are transported by laminar flow into the cavity. By solving the time-dependent convective-diffusion equation for the transport, the variation of the e.s.r. signal with time as the current is stepped on the electrode can be calculated. Theory and experiment are shown to be in good agreement. We have previ~uslyl-~ described how a tube electrode coupled to an e.s.r.spectrometer can be used to study the kinetics of radical species generated at the electrode. The electrode, which is a section of the wall of the tube, is situated just above the e.s.r. cavity. The species are carried by laminar flow from the electrode into the e.s.r. cavity where they are detected and their concentration measured from the e.s.r. signal. The variation of this signal with flow rate characterises whether the radical is stable,l whether it is decomposing by simple first-2 or ~econd-~ order kinetics or whether a more complicated kinetic scheme operate^.^ In this paper we consider the transient signal seen when the current at the electrode is changed. In particular, the case for a galvanostatic step is considered and theory and experiment are shown to be in good agreement.THEORY The convective-diffusion equation describing the concentration of a stable radical is where v, is the velocity of flow at the centre of the tube, ro is the radius of the tube, r is the radial distance from the centre of the tube, and x is the distance down the tube. The following variables are used to express eqn (1) in dimensionless variables t Present address: Department of Inorganic, Physical and Industrial Chemistry, The University of Liverpool, Liverpool L69 3BX. 51 1561 FAR 11562 TUBE ELECTRODE A N D E.S.R. TRANSIENT SIGNALS where xE is the length of the electrode and 1 is the total length of the electrode and cavity. Providing' that eqn (1) can be written (lD/uori) 4 1 av a2w az a t 2 ax (3) ---- - where the concentration c has been normalized with The relevant boundary conditions are at z = o , v = o t + c o , V+O t = O as and at for a t and for av a t The functionflz) describes the variation of the flux at the surface of the electrode XdXE, %=AT) x>&, -=o.A T ) = - We extend our previous function M3 to describe the transient e.s.r. signal S, normalised with respect to the current, i, at the electrode where vf is the volume flow rate (V, = xrEu0/2) and n is the number of electrons transferred in the electrode reaction. M , varies between 0 at z = 0 and M , the steady-state e.s.r. detection effi~iency,~ as z -, co. The e.s.r. signal S is given by1 where S* is the signal for one mole of spins at the centre of the cavity. Then from eqn (6) and (7) We now take the Laplace transform of eqn (3) with respect to time, and obtainW.J. ALBERY AND R. G . COMPTON 1563 where p is the transform variable. Eqn (9) has the same form as eqn (4) of ref. (2) which describes radical decomposition by first-order kinetics. The boundary conditions for the two equations become identical if c' in table 1 of ref. (2) is replaced byfip). Thus using the series solution previously obtained2 we can write where b, has been defined in ref. (2) and I 7=0 7 /- 0.5 1-0 1.5 2.0 I ... $ ... I I 1 I 7 FIG. 1.-Calculated e.s.r. signal transient due to a galvanostatic step caused by a step in the current at the electrode (as shown in the inset). So far we have not specified the current transient on the electrode. For the particular case of the galvanostatic step shown in fig. 1, f(z) = - 1 and s o f i p ) = - l/p.Hence (14) 51-21564 TUBE ELECTRODE AND E.S.R. TRANSIENT SIGNALS A power series such as eqn (14) cannot be inverted term by term. Consequently we adopt the method of rational approximation^.^ That is we take the first n terms of eqn (14), J,(p), and express them as a ratio of two polynomials of degrees n - 1 and n, which in turn yield n terms that can be inverted analytically. The procedure is repeated until increasing n makes no significant difference to the result. The particular procedure employed was Levin transformation6t7 of g ( p ) , as described in the Appendix. The function h(z) in eqn (12) is given by eqn (A 10). A complete programme was written to evaluate this function. Good convergence was obtained as n increased for all but small values of z (z c 0.3).For z = 0.2 no convergence was found even taking n as large as 17. Similar problems have been found with ring-disc transients. In contrast only a small number of terms were needed for larger values of z, for instance n = 5 for z > 0.8. Clearly the method of rational approximations is a powerful and efficient method of performing inverse Laplace transformation when analytical methods fail. In fig. 1 the converged values of h(z) are plotted to show a signal transient as a function of time. It is apparent that the transient reaches the steady-state value when z - 2.4. Assuming that D - 10+ cm2 s-l we find that this value of z corresponds to a real-time value of ca. 1 s for our apparatus at the fastest flow rate available.These times will be shorter for unstable species since their concentration profiles extend less far into the cavity. EXPERIMENTAL The basic apparatus and technique have been described previously.’. The aqueous sol- utions used contained 3 mmol dmP3 NNN”’-tetramethyl-p-phenylenediamine (TMPD) and 0.10 mol dm-3 K,SO,. The K,SO, was of AnalaR standard. TMPD was B.D.H. LR grade recrystallized as described by Michaelis.8 All solutions were made up in doubly distilled deionized water. The electrode was made of platinum and had the following geometrical parameters (uide supra), x,/cm = 0.29 and r,/cm = 0.045. All potentials were measured with respect to the saturated calomel electrode. The e.s.r. spectrometer used was a Bruker ER200tt.Transients were obtained by holding the magnetic field at the value corresponding to the maximum in the central peak of the spectrum. RESULTS AND DISCUSSION The system studied was the oxidation of TMPD: NMe2 NMe2 Steady-state studies on the tube electrode showed that the electron transfer was reversible and that the diffusion coefficient of TMPD was 4.9 x cm2 s-l. The e.s.r. spectrum obtained was in good agreement with the literature.g$ lo The spectrum was overmodulated in order to increase the sensitivity. Fig. 2 shows a log-log plot of the steady-state signal ( S / i ) against 5. The observed gradient of -2/3 is that expected from a stable radical.’ All the transients recorded at different flow rates were normalised onto one curve using eqn (2). The resulting plot is shown in fig.3. The solid line shown is theW. J. ALBERY A N D R. G. COMPTON 1565 1 o-2 to-' Vf/cm3 s-' FIG. 2.-Variation of the normalised e.s.r. signal (S/i) as a function of flow rate. The slope of -2/3 is that expected for a stable radical. f r-o-'-our /" 9' 7'' theoretical line derived above. Reasonable agreement is found between theory and experiment. Fig. 4 shows the time taken to reach half of the steady-state value as a function of flow rate. The theoretical line shown was calculated using the curve in fig. 1 and eqn (2), together with the experimental value for the diffusion coefficient. Again reasonable agreement is found between theory and experiment. Potentiostatic transients were also recorded using the same chemical system. The1566 TUBE ELECTRODE A N D E.S.R.TRANSIENT SIGNALS FIG. 4.-Variation with flow rate of the time taken for the signal to reach half of its steady-state value. The solid line is the theoretical line calculated using D = 4.9 x cm2 s-l. resulting signal transient was indistinguishable from galvanostatic transients measured at the same flow rate and having the same steady current value. This is because the large size of the cavity relative to the electrode means that although a larger number of radicals is produced during the potentiostatic step, this increase is small compared with the total number of radicals present in the cavity during the time that the transient is observed. The theory and experiments presented here therefore allow the e.s.r. tube electrode apparatus to be used for transient measurements.The advantage of such measurements over steady-state measurements is that surface processes such as adsorption can be studied. This approach is similar to that used in ring-disc electrode studies.ll9 l2 Using the observed transient and the transient calculated above it is possible to calculate the amount of radicals adsorbed at the electrode. For the existing tube e.s.r. apparatus one or more monolayers would be required for the method to be applicable. The severity of this restriction arises from the relatively long transit times between electrode and cavity. The sensitivity would be significantly increased if the electrode were placed at the centre of the cavity. The development of an in situ tube electrode will be reported in a forthcoming paper and this development should allow smaller quantities of adsorbed radicals to be studied.We thank the S.R.C. for financial support. APPENDIX This Appendix describes the inversion of g ( p ) in eqn (14) where and ck-l Jk-l a, =W. J. ALBERY AND R. G. COMPTON 1567 Each of the partial sums gn(p), where k=m gn(P> = x P”-’ k = l is transformed using the Levin U, transformation6$ ’ to yield the rational approximation U , where Un = Gn(P)/Hn(P) n i-1 G,(p) = E atpi-’ and The coefficients a, and pi can be found from the coefficients a,, a2, a3. . . , a,,, of the series (A 1) by means of the following equations6* (n + 2 ai= i ( ) J-1 J-l and On inverting eqn (A 4) we find where qm is given by (A 9) and Carrying out the integration in eqn (13) we may finally write for h(z) in eqn (12) H ~ ( P ) = d[Hn(p)I/d~- W. J. Albery, B. A. Coles and A. M. Couper, J. Electroanal. Chem., 1975, 65, 901. W. J. Albery, R. G. Compton, A. T. Chadwick, B. A. Coles and J. A. Lenkait, J. Chem. Soc., Faraday Trans. I , 1980, 76, 1391. W. J. Albery, A. T. Chadwick, B. A. Coles and N. A. Hampson, J. Electroanal. Chem., 1977,75,229. W. J. Albery, R. G. Compton and I. S. Kerr, J. Chem. Soc., Perkin Trans. 2, to be published. Y. L. Luke, Q. J. Mech. Appl. Math., 1964, 17, 91. D. Levin, Znt. J. Comput. Math., 1973, 3, 371. I. M. Longman, SZAM J. Appl. Math., 1973, 24, 429. * L. Michaelis, M. P. Schubert and S. Granick, J. Am. Chem. Soc., 1939, 61, 1981. J. R. Bolton, A. Carrington and J. des Santos-Veiga, Mol. Phys., 1962, 5, 615. lo K. H. Hauser, Mol. Phys., 1963, 7 , 195. l1 W. J. Albery, R. G. Compton and A. R. Hillman, J. Chem. Soc., Faraduy Trans. I , 1978, 74, 1007. W. J. Albery and A. R. Hillman, J. Chem. SOC., Faraday Trans. I , 1979, 75, 1623. (PAPER 1 / 1069)

ISSN:0300-9599    

DOI:10.1039/F19827801561

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J . Chem. SOC., Faraday Trans. I , 1982, 78, 1569-1578 Solvent Isotope Effects on the Enolisation of Acetone BY W. JOHN ALBERY* AND JONATHAN S. GELLES Department of Chemistry, Imperial College of Science and Technology, South Kensington, London SW7 2AY Received 7th July, 1981 The solvent isotope effect on the enolisation of acetone has been measured in acetic acid/acetate buffers. The precision of the differential method has been further improved by including differential dilution. Results have been obtained in both D,O and in a 50: 50 mixture of H,O and D,O. The number of sites and their fractionation factors have been found for the six different transition states including the concerted mechanism involving HAc and Ac-. The structure of the different transition states is discussed.The prototropic rearrangement by which acetone is converted to its enol has been much studied.’ The reaction is interesting in that there are two proton transfers, one in which the C-H proton is transferred to a base B and the other in which the carbonyl oxygen is protonated by an acid HA. The reaction can therefore be written as in the following scheme: (Me),CO + (HA + B) or (A + HB) 1 [B - - -H - - --CH,C(Me)CO---- H ----A]$ 4 C 1 40 CH, : C(Me)OH + (HA + B) or (A + HB). In this scheme dC and #o are the fractionation factors, which describe the effects in the transition state of substituting D for H for the two protons that are transferred from carbon and to the enol oxygen, respectively. In an acetic acid/acetate (HAc,/Ac-) buffer the velocity of the reaction, v, obeys the following rate law2: The second form of the equation shows how the different terms depend on the buffer constituents.In this paper we report data for the solvent isotope effect in HDO and D,O for each of the terms in eqn (2). Combination of these data with previous results 15691570 SOLVENT ISOTOPE EFFECTS for C-H isotope effects (&) and for Brarnsted slopes allows us to determine the positions of the protons marked by #c and #o in eqn (1) in each of the six transition states corresponding to the terms in eqn (2). In particular we are interested in whether the protons move in a stepwise or in a synchronous fashion. A simple model, based on the Marcus t h e ~ r y , ~ . ~ which describes the results, is presented in the subsequent paper.THE DIFFERENTIAL METHOD More information about transition states can be obtained by measuring the solvent isotope effect in an equimolar mixture of H,O and D,O (HDO) as well as in D,O itself. But to be successful the data must be very precise. We have shown previously5 that the differential method can be used to study the iodination of acetone. In this method the difference in reaction rate between, say, H,O and HDO is measured directly by using a double-beam spectrophotometer to monitor the iodine concentration simultaneously in both solutions. This technique leads to greater precision in the kinetic data, but to obtain any advantage one needs an equal precision in knowledge of the constituents of the two different solutions. To achieve this we have modified the method. A volume, Vk, of a stock solution of buffer and acetone in H,O is diluted by weight with an approximately equal volume, VDz0, of D,O to provide the HDO reaction mixture.A second volume, &, of the same stock solution is diluted by weight with a volume, VHzo, of H,O to provide the H,O reaction mixture. Small quantities of iodine are then added to the two solutions in the double-beam spectrophotometer. The reactions are zero order in iodine. If gHBO is the gradient of the H,O optical density with time and 6g the difference in gradient (gHzO -gH,o), which is observed directly by the spectrophotometer, then the isotope effect is given by: where n is an exponent depending on the order of the reaction for the particular term in eqn (2) being studied.The volume 6 for the H,O dilution is chosen so that 6g is as small as possible. In the ideal experiment no change in optical density would be observed and the solvent isotope effect would be simply determined by the two weighed dilutions of the stock solution. A typical trace is shown in fig. 1. In order that gHzO can be observed, approximately twice as much iodine is added to the H,O mixture. By starting with a D,O stock solution and diluting with D,O and H,O (to HDO) the rate ratio kDBo/kHDo can be measured in a similar manner. EXPERIMENTAL The experimental details of the differential method have been described previ~usly.~ All reactant solutions were mixed in flasks in a thermostat tank and transferred to the Cary 14 spectrophotometer using thermostatted pipettes.Thermocouple measurements showed that the temperature difference between the two cells in the spectrophotometer was < 10-30C. Swapping the positions of the two reactant solutions made no significant difference to the results. The reactions were followed by measuring the zero-order removal of I, by monitoring the concentration of I; at the wavelength of its maximum absorption, i/nm = 351. The initial concentrations of I- and I, were ca. 30 mmol dmP3 and 30 pmol dmV3, respectively. About twice as much I, solution was added to the solvent containing more H so that g could be observed.W. J. ALBERY A N D J. S. GELLES 1571 - time FIG. 1 .-Typical experimental trace of optical density against time for a double differential experiment. The broken lines show the variation of optical density in each of the two cells.The gradients of optical density with time for the end of the H,O reaction, gH,O, and for the earlier differential trace, dg, are measured. The dilutions are arranged so that the rates of consumption of I; are almost the same in the two cells, and dg is therefore close to zero. These additions were made from a syringe. In each case the volume added was so small that it did not affect the dilution ratio in eqn (3) or the temperature. All chemicals were AnalaR grade and all water was doubly distilled. Buffer solutions were made up by titrating acetic acid with sodium hydroxide using a Radiometer TTT 2 titrator. All D20 was distilled under N, to remove a significant quantity of some basic impurity. The isotopic content of H 2 0 + D20 mixtures was determined either by adding a known quantity of methanol and measuring the n.m.r. spectrum or by the infrared method of Kreevoy and Straub.6 In making up reaction solutions the time of contact of acetone with deuterated solvents was minimised so that insignificant isotope exchange took place on the acetone.RESULTS AND DISCUSSION Results for catalysis by H30+ were obtained in solutions containing between 2 and 3 mmol dm-3 HClO, and ca. 0.1 mol dm-3 acetone. Results for the other catalysts were obtained in buffers of acetic acid. The conditions were chosen so that each of TABLE VA VALUES OF RATE CONSTANTS IN EQN (2) (IN mol dm-3 s-l) catalyst Bell2 this work H,O+ 2.7 x 10-5 2.8 x 10-5 OH- 0.25 0.25 H2O 8 x 8 x HAc 8 x lop8 7 x 10-8 Ac- 2.5 x 10-7 i.9x 10-7 HAc, A c - ~ 3 x 10-7 2 x 10-7 a Units are mo12 dm-6 s-l.the terms in eqn (2) was dominant. Fig. 2 shows the zone of dominance of each of the terms and the composition of the buffer mixtures used in our experiments. From the values of gHgO, the gradient of optical density against time, observed at the end of the differential experiments (see fig. l), we find the values for the individual1572 SOLVENT ISOTOPE EFFECTS I 1 .o FIG. 2.-Contour diagram showing the largest term in eqn (2). The zone of dominance of each term is shown by the full lines. The broken lines are contours showing how the observed rate constant vanes with the buffer constituents. The asterisks show the buffer concentrations used in this work. rate constants in eqn (2) given in table 1.Good agreement is found with the previous results of Bell and coworkers.2 The differential method with differential dilution was tested by carrying out H30+ catalysed reactions with H 2 0 in both cells. The mean of 5 determinations was 1 .ooo & 0100 1. In the measurement of the solvent isotope effects the atom fraction of deuterium was usually a little less than 0.50 and 1.00. Small linear corrections were made as follows : (4) and where kI)20/kH20 is found by iteration. A set of typical results is given in table 2.W. J. ALBERY AND J. S . GELLES 1573 TABLE 2.-TYPICAL RESULTS FOR CATALYSIS BY L30+ XHDO 0.485 0.486 0.490 0.491 0.492 0.493 0.493 0.495 1.412 1.393 1.437 1.410 1.397 1.416 1.423 1.420 1.430 1.409 1.449 1.421 1.407 1.425 1.432 1.426 - 1.425 - f 0.005 - - a Calculated from eqn (3); calculated from eqn (4) with kDzO/kHzO = 2.18.TABLE 3.-RESULTS FOR THE SOLVENT ISOTOPE EFFECT IN ACETIC ACID BUFFERSa catalyst OH- H2O HAc Ac- HAc Ac- 5 x 10-3 6 x lop3 0.87 0.96 0.771 0.003 0.030 0.033 0.65 0.71 0.728 0.002 0.26 0.26 0.38 0.38 0.717 0.014 0.5 1 0.49 0.97 0.95 0.750 0.01 3 1030 1210 25.5 29.9 0.825 0.002 1020 1110 33.7 36.6 0.819 0.007 ____ ~~ 3.59 1465 3.72 1469 454 495 467 497 0.897 0.84 I 0.005 0.012 1.57 1.66 43 6 446 1 0.915 0.006 a All concentrations are in units of mmol dm-3. Results from ten experiments comparing catalysis by L,O+ in HDO and D,O gave kD2,/kHD0 = 1.533 0.005. (6) The acetic acid/acetate buffer concentrations and the results for the solvent isotope effect, after correction with either eqn (4) or eqn (5) are given in table 3.While reaction conditions are chosen so that one term in eqn (2) is as dominant as possible, the other terms cannot be completely neglected (see fig. 2). We write for each dominant term, DT, and1574 SOLVENT ISOTOPE EFFECTS Each 6 is a correcting ratio, for say kx, of the form and px and qx are the orders of the X term in eqn (2) with respect to HAc and Ac-, respectively. In eqn (9) and (10) we have introd7Lt:ed y+ and y , which describe the isotopic fractionation in the transition state for each catalytic route. The terms in 4 H A c and QAc- allow for the reactant fractionation. From analysis of thermodynamic' and kinetics data plausible values for 4HAc and QAc- are9 4 H A c = 0.98 and QAC- - 0.90.We use these values, but we have also repeated the calculations using &Ac = @,,- = 1 .OO and the conclusions are not affected. Values of the solvent isotope effects for each route and of the transition state fractionation described by yi and y,, where were calculated using eqn (7)-( 11) and are given in table 4. The values for L,O+ have been calculated from the results in table 2 and eqn (6) usinglO7 l1 I = 0.69 and1, pL+ = 0.03, where pL+ is a correction for the breakdown of the role of the geometric mean. are similar. The product y,, which is equal to II4$, is compared with values calculated from previous work and found to be in reasonable agreement. In our work the fractionation in the OH- transition state has been measured without using OH- as a reactant and hence we obtain our result without having to make any assumptions about OH- reactant fractionation.Interesting information about the structure of the transition state can be found not only from the overall size of the solvent isotope effect (kDZO/kHz0) but also from the way that the solvent isotope effect varies with the deuterium content, x, of H,O + D,O mixtures. For instance if a single site was responsible for the solvent isotope effect then k,/kHZo would be a linear function of x. On the other hand if general solvation on four or more sites was responsible then 1n(k,/kH2,) would be a linear function of x. In most cases the difference between these two extreme cases is small and that is why data of high precision are required. One finds the maximum difference when in HDO with x = 1.We have shown that no more information can be found by studies at other values of x. Hence at best only one extra piece of information can be obtained by the study of solvent isotope effects at intermediate values of x and there is no way in which one can obtain a 'proton inventory'. The extra piece of information is best expressed as the curvature parameter, y,13 where y describes the curvature in a plot of ln(k,/k,20) against x. If this plot is a straight line (showing four or more sites) then y = 0. At the other extreme with only one site, the In plot is curved and this It can be seen that for each catalyst as expected the results for yi andW. J. ALBERY A N D J. S. GELLES 1575 TABLE 4.-RESULTS FOR TRANSITION STATE FRACTIONATION catalyst H,O+ OH- H2O HAc Ac- HAc Ac- kHDO/kHzOd 1'425 Y? 0.853 kDnO/kHDOu 1 S30 Y? 0.839 YIC 0.716 Y l Yh 0.6 n' 2 (Pck 0.13f 0.1 3m 0.15P +/- 0.2 0.77 1 0.741 0.728 0.707 0.524 0.60" 0.4 0.1 2 or 3 0.10" 0.1 3n 0.14 0.12 0.717 0.56 0.750 0.53 0.30 0.2 0.3 3 2 0.825 0.860 0.819 0.820 0.706 0.63f 1.6 0.2 1 0.12f ~~ ~ 0.897 0.841 0.886 0.72 0.915 0.926 - 0.820 - 0.77f 0.449 - 1 .o 0.2 - - ? 1 0.21f 0.179 a Observed rate ratio from double differential method after correction from eqn (4) or (5).Transition state fractionation calculated from eqn (7)-( 1 1) assuming for reactant fractionation that (PLAc = 0.98, OAc- = 0.90, 1 = 0.69 and pL+ = 0.03.12 0. Reitz, Z . Phys. Chem., Teil A , 1937,179, 119. J. R. Jones, Trans. Faraduy SOC., 1965,95,61.f 0. Reitz and J. Kopp, 2. Phys. Chem., Teil A , 1939, 184, 429. A. F. Hegarty and W. P. Jencks, J. Am. Chem. SOC., 1975, 97, 7188. Curvature parameter13 calculated from eqn (12). J n 3 y-l. 0. Reitz, Z . Elektrochem., 1937, 43, 659. Y. Pocker, Chem. Ind., 1959, 1383. P J. Toullec and J. E. Dubois, J. Am. Chem. SOC., 1974, 96, 3524. y, = y i x y ; . Fractionation factor for L in flight from C assuming secondary factors are unity. curvature gives a value of y = 1.00. From the values of yi and y! we can calculate values of y, where13 (12) 4(lnyi - lnyd) ' = (lnyg + lnyd)2 * A knowledge of y is useful because we have shown13 that the number of sites contributing to the solvent isotope effect must be greater than or equal to n, where n = y - l . (13) In table 4 we report values of y and n.Finally we report values for the fractionation factors, &, for the proton or deuteron being transferred from carbon. The values for these factors (0.12-0.20) are all much smaller than the values of y , (0.30-0.82). Now y , is a product of factors most of which are less than unity; hence the factor for the proton being transferred to oxygen, &, must be greater than y, and hence for each transition state significantly greater than &. In each transition state the proton being transferred from carbon is in flight while the proton being transferred to oxygen is at least partially bound. The Brarnsted slopes for catalysis by a series of carboxylic acids,14 carboxylate bases14 and for the third-order push-pull route15 are a = 0.55, p = 0.88 and p- a = 0.17, respectively.Bell and Jones16 have shown that because a (for the acids) plus a (for the bases) does not equal unity, the transition states for catalysis by acids and by bases must be different. Following these arguments and those of Swain and c o ~ o r k e r s ~ ~ ~ l8 and of Lienhard and Wanglg we locate the transition states for acid1576 SOLVENT ISOTOPE EFFECTS ,Me Me H-CH, -C* CH2 = CCoH OH' ACID HAc H,O' HA * c IAc- +/ ,Me M e H-CH2-C - - CH2=C', +O 0- FIG. 3.-A map showing the location of the six transition states corresponding to the six terms in eqn (2). and base catalysis as shown in fig. 3. For the acid-catalysed route the observed a arises from a contribution of unity for the protonation of the carbonyl group and a value ofg = 0.45 for the subsequent base catalysed removal of the C-L proton (a = 1 -p).As argued by Hegarty and Jencks15 the transition state for the push-pull route must lie near the centre of the diagram. There is compensation between the push and the pull as KA is changed resulting in a small value for the Brransted slope. With these considerations in mind we display in fig. 3 plausible structures and fractionation for the six transition states in table 4. For the H30+ transition state from the Hammond postulate20 or the Marcus t h e ~ r y , ~ since H,O is a weaker base than Ac-, /? must be greater or equal to the value of 0.45 for the carboxylic acids. If we attribute, in accordance with the value of y, all of the observed solvent isotope effect to the L,O molecule receiving the flying C proton, then from the otyerved value of y 1 = 0.72 we deduce a fractionation factor for each solvent site of yi = 0.85 and from the Gold-Kresge relation2l7 22 = In (0.85)/1n (0.69) = 0.45. (14) Hence the fractionation factor for the enol site must be close to unity. If a factor less than unity is attributed to this site then PI < 0.45, violating the Hammond postulate.Next we note that the values of 4c for the symmetrical H,O+ and HAc transition states are small and similar (0.14 and 0.12, respectively). These values may be contrasted with that for the Ac- transition state where dC = 0.21. This transition state is very product-like ( B = 0.88) and hence on the Westheimer principle23 we expect the primary factor to be closer to unity.The solvent isotope effect arises from the developing solvation of the enolate anion. For the OH- transition state p must be significantly less than the value for the carboxylate bases, because, first bc (= 0.12) is again close to the symmetrical values rather than the value for Ac-, and secondly y1 is much smaller for OH- than for Ac- (0.52 compared with 0.8). The lower value of y1 must be caused by significant lyoxide fractionation. We assume for OH- in the transition state that one lone pair on the oxygen is taken by the proton flying from carbon leaving two solvent protons hydrogen-bonded to the remaining lone pairs, and one solvent derived proton as LO-. A rough estimate from the known parameters2* for OH- suggests that /? z 0.5. On changing from Ac- to OH- the transition state becomes significantly more reactant-like.W.J. ALBERY AND J. S . GELLES 1577 A shift in the opposite direction takes place on changing from Ac to H20. The low solvent isotope effect (’yl z 0.3) arises from the almost complete development of L,O+ and enolate in the very product-like transition state. Here we have assumed that L 2 0 reacts by the base catalysed route. The alternative acid catalysed route, in which there is auto-protolysis to L,O+ and OL-, is considered in the subsequent paper. With regards to the push-pull mechanism the value of y of unity supports the conclusion of Hegarty and Jencks15 that this transition state is close to the centre of the diagram and is the one transition state which has the enol proton almost in flight, with Qo = 0.5.Factors of this size have been found by Kreevoy and for L when it is in a stable hydrogen bond between two bases of similar pKvalues. Hence in the push-pull transition state, while the carbon proton is flying through the air (Qc = 0.17), the oxygen proton is transferred in a stable hydrogen bond. This type of transfer was suggested by Swain et a1.26 and by Kreevoy and Corde~.~’ As the carbon proton is removed, the base strength of the carbonyl oxygen increases, and when it is equal to that of the catalysing acid HA the proton transfers within the W-shaped well. The value of Qc (0.17) may be closer to unity than for the other symmetrical transition states, because the reaction co-ordinate contains some significant motion from the enol proton. Finally we come to the puzzle of the HAc transition state.The carbon proton is roughly half transferred to the acetate ion. There is significant solvent fractionation on a single site which must be the enol proton with Qo z 0.75. But this value seems to contradict the value of & = 1.00 deduced above for the very similar H,O+ transition state. This significant difference arises from the fact that the values of y1 for the two transition states are very similar (0.716 and 0.706) but the H,O+ transition state has the two sites on L 2 0 receiving the carbon proton whereas the HAc transition state only has Ac- solvation of @I,-. The reason for this difference will be explored in the subsequent paper, where the Marcus theory will be applied to provide a quantitative explanation of the six different transition state structures suggested in this paper.We thank Prof. Kreevoy for helpfui conversations and the S.R.C. both for a studentship for J.G. and for the purchase of a spectrophotometer. R. P. Bell, The Proton in Chemistry (Chapman and Hall, London, 1973), pp. 141 et sey. R. P. Bell, The Proton in Chemistry (Chapman and Hall, London, 1973), p. 150. R. A. Marcus, J . Phys. Chem., 1968, 72, 891. R. A. Marcus, J . Am. Chem. SOC., 1969, 91, 7224. W. J. Albery and B. H. Robinson, Trans. Faraday SOC., 1969, 65, 980. M. M. Kreevoy and T. S . Straub, Anal. Chem., 1969, 41, 214. V. Gold and B. M. Lowe, J . Chem. SOC., A, 1968, 1923. W. J. Albery, A. N. Campbell-Crawford and R. W. Stevenson, J . Chem. SOC., Perkin Trans. 2, 1972, 2198. W. J. Albery, in Proton-Transfer Reactions, ed. E. F. Caldin and V. Gold (Chapman and Hall, London 1975), p. 285. l o A. J. Kresge and A. L. Allred, J . Am. Chem. SOC., 1963, 85, 1541. l 1 V. Gold, Proc. Chem. Soc., 1963, 141. l2 W. J. Albery, in Proton-Transfer Reactions, ed. E. F. Caldin and V. Gold (Chapman and Hall, London, 1975), p. 280. l 3 W. J. Albery, in Proton-Transfer Reactions, ed. E. F. Caldin and V. Gold (Chapman and Hall, London, 1975), p. 273. l 4 R. P. Bell and 0. M. Lidwell, Proc. R. SOC. London, Ser. A , 1940, 176, 88. A. F. Hegarty and W. P. Jencks, J. Am. Chem. SOC., 1975, 97, 7188. l6 R. P. Bell and P. Jones. J . Chem. SOC., 1953, 88. C. G. Swain, A. J. DiMilo and J. P. Gardner, J . Am. Chem. SOC., 1958, 80, 5983. C. G. Swain and A. S. Rosenberg, J . Am. Chem. Soc., 1961, 83, 2154.1578 SOLVENT ISOTOPE EFFECTS l 9 G. E. Lienhard and T. G. Wang, J. Am. Chem. Soc., 1969, 91, 1146. *O G. S. Hammond, J . Am. Chem. SOC., 1955, 77, 334. 21 V. Gold, Trans. Faraday SOC., 1960, 56, 255. 22 A. J. Kresge, Pure Appl. Chem., 1964, 8, 243. 23 F. H. Westheimer, Chem. Rev., 1961, 61, 265. 24 V. Gold and S. Grist, J. Chem. Soc., Perkin Trans. 2, 1972, 89. 25 M. M. Kreevoy, T. M. Liang and K. C. Chang, J . Am. Chem. SOC., 1977, 99, 5207. 26 C. G. Swain, D. A. Kuhn and R. L. Schowen, J. Am. Chem. SOC., 1965, 87, 1553. 27 E. H. Cordes, Prog. Phys. Org. Chem., 1967, 4, 1. (PAPER 1 / 1078)

ISSN:0300-9599    

DOI:10.1039/F19827801569

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J . Chem. SOC., Faraday Trans. 1, 1982, 78, 1579-1590 Application of the Marcus Relation to Concerted Proton Transfers BY W. JOHN ALBERY Department of Chemistry, Imperial College of Science and Technology, South Kensington, London SW7 2AY Received 7th July, 198 1 A model, based on the Marcus theory, is presented for the ket-no1 tautomerism of acetone. With one adjustable kinetic parameter, the model explains the rate constants and isotope effects for the six different catalytic routes involving H,O+, OH-, H,O, CH,COOH, CH,COO- and the third-order term with both CH,COOH and CH,COO-. The range of pK, over which the third-order term is observed is also explained. General conclusions are derived and discussed concerning the catalytic advantage of the concerted mechanism described by the third-order term.In the previous paper1 we presented data for the solvent isotope effect on the kinetics of the enolisation of acetone. The reaction may be written: (Me),CO + (HA + B) or (A + HB) 1 dc 1 4* Ah (1) [B-- --H _ _ _ -CH,C(Me)O _ _ _ _ H - _ _ _ CH,: C(Me)OH + (HA + B) or (A + HB). In acetic acid/acetate buffers we have to consider the nine different transition states listed in table I , which can give rise to the six different catalytic rate constants discussed in the previous paper.l From the isotope effects and the Brsnsted slopes we deduced structures for the six different transition states. In this paper we apply the Marcus theory to the transfer from carbon and show that with only one adjustable parameter we can obtain a reasonable explanation for the free energies of the six different transition states, the isotope effects and the Brsnsted slopes.In addition we show that the three missing transition states in table 1 have higher free energy than those that are in fact observed. THE MODEL In eqn (1) one of the protons is transferred from carbon while the other is transferred from an oxygen base to the carbonyl oxygen. The fractionation factors for the carbon proton are all small1 (0.12-0.21) and show that this proton is in flight. By contrast the proton being transferred to oxygen is always at least partially bonded. Hence we assume that the main activation process is concerned with transferring the carbon proton and that negligible activation energy is involved in transferring the proton to oxygen.We further assume that the proton transfer from carbon is governed by the Marcus theory,,? but the thermodynamics for this transfer depends on the degree of proton transfer to the carbonyl oxygen. If there is no transfer then the reaction (Me),CO+B- -+ CH,:C(Me)O-+HB (1’) 15791580 MARCUS THEORY APPLIED T O PROTON TRANSFERS TABLE 1 .-NINE TRANSITION STATES FOR EQN (1) rate $ HA B constant CPKCU A P P I 2 (3 4 ( 5 6 (7 8 9 H,O+ H,O+ HAC H,O+ HAc HAc H2O H2O H2O 1 1 0 17.4)" 0 1 - 1" 1 6.4' 1 1 1 .O)' 1 2 - - - - - Free energy involved in forming caged complex from reactants. Free energy involved in forming HA and B from A and HB when B is a stronger base than A; ApK = pKLB - p&+. ' This transition state is of higher free energy than the one following it in the table and so is not observed.I 1 I 1 A H ~ = C - C - H B A HO-C=C HB HA, A I BASE I I I I I AH -0-C-C HB 0- A H 0.C-C-H B I FIG. 1 .-Location of the transition states for the base-catalysed, concerted and acid-catalysed routes for the enolisation of acetone. is very uphill. On the other hand if this proton is transferred, the reaction (Me),CO+H + B- -+ CH,: C(Me)OH + HB may well be downhill. Fig. 1 is helpful in showing different possible transition states, where the degree of proton transfer from carbon is given by p and the degree of proton transfer to oxygen by a. We assume that for intermediate values of a the thermodynamics for the Marcus transfer from carbon is given by a linear free energy relationship between the two extremes of a = 0 and a = 1.This type of model has been discussed for proton transfers to diazo corn pound^.^ Finally we assume that theW. J . ALBERY 1581 intrinsic kinetic barrier for the exchange at carbon does not depend on a. The size of this barrier is then the only adjustable kinetic parameter in the model. Before calculating the kinetic barriers, we have first however to discuss the thermodynamics which fix the free energy differences between the four corners of fig. 1 . Each of these corners represent the caged reactants, intermediates or products. These caged encounter complexes are formed from free reactants. In those cases where B is a stronger base than A then free energy, corresponding to ApK, must be expended in transferring a proton from B to A; values of ApK are given in table 1.Next we assume that the association of a catalyst with the substrate has an association (2) constant, K,, given by Kc/dm3 mol-l = 0.1. Values of CpK, are given in table 1 ; the water-catalysed reaction has CpK, = 0 since the water is always present, and the third-order term has CpK, = 2 since two molecules of catalyst are involved. Each catalytic route k, is then described by a first-order rate (3) constant given by where n is the order of the observed rate constant. With all the rate constants reduced to first-order rate constants, to describe the inter-conversion of the species depicted in fig. 1 it is convenient to express the acid-base equilibria in dimensionless equilibrium constants. We do this by considering the equilibrium k;, = (Mobs/K:-l HA + H,O $ H,O+ +A- Ka.Introduction of the concentration of H20, (55.5 mol drn-,), gives where pKHA is the more usual scale, in which KHA has the units of mol drn-,. enol content of acetone is ca. lo-’ and that for the equilibrium As regards the values for acetone, Guthrie and c o ~ o r k e r s ~ - ~ have shown that the (Me),CO+H,O~CH,:C(Me)O-+H,O+ Kkz0 the value of (PK’)~=~ is 21. There is more controversy concerning the protonation of acetone which fixes the fourth corner: (Me),CO+H + H 2 0 Me2C0 + H,O+ KiH+. E~tirnates8-l~ of pKs,+ range from - 2.2 to - 7.2. The isotopic data presented in the previous paper show that the H,O+-catalysed transition state has a < 0.5. However, values of pKsH+ of -2, implying values of pK&,+ of about zero, would mean that the transfer would be uphill with D > 0.5.For this reason we consider that pKs,+ must be at the lower end of the range and we take pK&,+ = -6. This means that in eqn (5) (Me),CO+H+H,O + H,C:C(Me)OH+H,O+ KkZ1 ( 5 ) pKk=, = 1. THE MARCUS EXPRESSION The Marcus expression is normally written in terms of free energies, but in dealing with proton transfers, where pK values are more familiar quantities, it is more instructive to write the expression using pk’ where pk’ = -log (k’h/kT) = AGJ/2.3 RT1582 MARCUS THEORY APPLIED TO PROTON TRANSFERS and AGJ refers to the first-order rate constant. In using the values of K, in eqn (2) and table (1) we have already dealt with the problem of wR. For any value of a in fig. 1 the Marcus expression gives pkh = ipk; ++ApK; +(A~K;)~/8pk;.(6) In this expression pkh refers to the rate for removing the carbon proton at the particular value of a. To this contribution must be added, first that concerned with reaching that value of a at p = 0 from the reactant corner and secondly that arising from the creation of the catalysts HA and B for those cases where B is a stronger base than A: Returning to eqn (6), pk; is the term corresponding to the degenerate reaction for Me,CO + H2C: C(Me)OH s H2C: C(Me)OH + Me,CO. acetone : There is no similar term for the catalyst, bemuse we assume that the barriers for the degenerate catalyst reactions are much smaller than that for acetone. The term ApK; describes the thermodynamics for removing the carbon proton at a, and we assume that it is given by a linear free energy relation: pk’ = pk’, + a(pK;IA - pK&+) + ApK.(7) APK; = apK:,=l+(l-a)pK,:,,-pK;,B (8) where for acetone and pK;=, = 21. pK;=, = 1 THREE TYPES OF TRANSITION STATE From the isotope effects for the carbon proton, where & < 0.21, that proton is always in flight in the transition state and this implies that p has a fractional value between 0 and 1. There are then three types of transition state depending on the value a = 0 ‘basic’ of a : a = 1 ‘acidic’. 0 < a < 1 ‘concerted’ Fig. 2-4 show typical pk surfaces for the three different types of transition state. From eqn (6)-(8) we can write down expressions for the rate constant, for p and in the case of the concerted case for a. The values of a and p for the concerted case are given by differentiation to find the transition state of lowest free energy.The results are: basic a=O acidic a = lI 0 1 U 0 II U cd *-’ E d .r. v) .-1584 MARCUS THEORY APPLIED TO PROTON TRANSFERS C- OH c =o HAc t C =OH FIG. 4.-Typical pk' surface for the acid-catalysed route; the free energy required to form H,O+, Ac- from HAc is also shown, Note that this figure is viewed from the opposite side to those in fig. 2 and 3. The parameter pKA describes the following equilibrium : CH,C(Me)CO + CH,=C(Me)OH + CH,C(Me)CO+H + CH,=C(Me)CO-. This equilibrium is independent of the catalysts. It is a property of the substrate and is concerned with the free energies of the two unstable corners of the square (NW and SE) compared to the two stable corners (SW and NE).LOCATION OF THE TRANSITION STATES We start by considering a particular base B. Then the values of /?a=o for the base-catalysed reaction by B and of Pa = for the acid-catalysed reaction by HB are given by eqn (10) and (12), respectively. Between these two extremes there will be a range of concerted transition states whose locations lie on the line given by eqn (1 5). The locus of this line depends on the substrate parameters (pk,, pKI, = and pKA) and on pKhB but not on pKhA. Hence there is a line for each base, B. The concerted transition states are only found for a limited range of acids HA with pK' values which lie between pKLA, I and pK;IA, I I where The separation between pK&A, I and pK;IA, I I does not depend on the base B. For the concerted transition states from eqn (14) the value of /? depends on the substrate parameters pKLH and pKA and on PKLA but not on pK;IB.Hence each concerted transition state is located by pKLB on a line given by eqn (15) and its value of /? is given by pKhA and eqn (14). This pattern is illustrated in fig. 5.W. J . ALBERY I 1 A HO+=C-C-H B A HO-L=b=HB 1585 AH - O - b = c HB I 1 I 1 AH O = C - C - H f3 FIG. 5.-Pattern underlying the location of the transition states. Each acid, e.g. HA, is associated with a vertical line given by eqn (14). Each base, e.g. B, or B,, is associated with a slanting line given by eqn (1 5). Possible transition states lie at the intersections of the ‘base lines’ with either the ‘acid lines’ or the edges of the diagram. THE ACETONE SYSTEM We now apply the equations introduced above to the acetone system.In the model we have assumed that there is only one kinetic parameter associated with the substrate, pk;. This assumption is almost certainly an oversimplification, but, as shown in fig. 6, a reasonable fit between the observed and calculated pk’ values for the six acetone transition states can be found using pki = 19. We can also calculate pk’ values for the other three transition states in table 1. In each case the pk’ value of these three transition states is significantly higher than the observed transition state which corresponds to the same term in the rate equation. We can then calculate the locations of the six transition states using eqn (lo), (1 2), (14) and (15). These are plotted in fig. 7.The value of p = 0.45 for the acetic acid transition state is in excellent agreement with the slope of the Brsnsted plot14 for catalysis by carboxylic acids. Rather less good agreement is found for the acetate transition state with = 0.70 as compared with a Brsnsted slope of 0.88.14 The locations of the H,O+OH- and H,O transition states agree well with the locations deduced from the solvent isotope effects in the previous paper.l Assuming that the fully developed enolate ion has fractionation of 0.50 we obtain the results in table 2. For the case of OH- we have assumed that one of the oxygen lone pairs is receiving the flying proton from carbon, leaving two solvent protons hydrogen bonded to the remaining lone pairs and one solvent derived proton, LO-. Returning to fig.7 the location of the concerted HAc/Ac- transition state also agrees with the previous arguments.l? l5 In fig. 8 we show the location of the concerted transition states that arise from different conjugate acid-base pairs HA, A, together with the pK values of the catalysing acid. The calculated values of the rate constants hardly alter as the catalyst changes. This agrees with the very low slope of the Brsnsted plot found by Hegarty1586 MARCUS THEORY APPLIED TO PROTON TRANSFERS 25- 20- FIG. 6.-Comparison of observed and calculated values of pk' for pkl, = 19. The points shown by crosses are calculated values for three alternative transition states which are not observed. HB HB FIG. 7.-Location of the transition states. The brackets indicate the alternative transition states which are not observed.W.J. ALBERY 1587 TABLE 2.-sOLVENT ISOTOPE EFFECTS OH- 0.55 0.54 0.56 H2O 0.78 0.33 0.30 A H;=~-+.-H B A HO-b=C HB AH 0-C=C HE 0- A H O=b-b-H B FIG. 8.-Location of the concerted transition states for different conjugate acid-base pairs HA, A. The numbers indicate the pK of the acid and the solid line the range of pK for which the third-order term has in fact been observed.16 Because of the competition from other routes it is difficult to observe the third-order term near the edges of the diagram. and Jencks.16 Furthermore the range of catalyst pK over which Hegarty and Jencks observed the third-order term agrees almost exactly with that found from the model. Further support for the model arises from the variation of & !he fractionation factor for the carbon proton.The results in table 3 show that & is lowest for the symmetrical transition states and is closer to unity as increases. This pattern is that predicted by Westheimer for primary kinetic isotope effects.16 The only remaining isotope effect that is curious is the factor, do, for the oxygen site for catalysis by HAc. Table 4 collects together the data on q50. The local environment of the oxygen proton will depend on the relative acidities of the protonating acid HA and the changing basicity of the carbonyl oxygen, pKb,, where we can write the linear free energy relation as In this equation pKA is defined in eqn (16) and is equal to 20 for acetone. The equation describes how the basicity of the carbonyl oxygen varies from a pK of - 6 for p = 0 to a pK of + 14 for the enol ( p = 1).Values of pKbH and the difference in pK around the oxygen proton are also given in table 4. It is interesting that for every concerted transition state from eqn (14) and (19) we find the simple result that ApK' = 0. In the transition state the base strengths on either side of the oxygen proton are in balance. This agrees with our previous conclusions.4 For the HAc/Ac- case the fractionation1588 MARCUS THEORY APPLIED TO PROTON TRANSFERS TABLE 3.-vARIATION OF 4~ WITH p catalyst D 4c HAc 0.45 0.12 H30+ 0.52 0.14 OH- 0.55 0.1 1 HAc/ Ac- 0.62 0.17 Ac- 0.70 0.2 1 TABLE 4.-VALUES FOR $0 AND P K ~ H catalyst 40 D PKOH P K H A ApK H30+ 1 .o 0.55 4.2 0.0 4.2 HAc 0.71 0.43 2.6 0.0 2.6 HAc/ Ac- 0.51 0.62 6.4 6.4 0.0 factor is similar to that found by Kreevoy et al.” for symmetrical hydrogen bonds.By contrast for the H,O+ transition state the proton is firmly on the carbonyl oxygen with a fractionation factor close to unity. For the HAc transition state the proton is still on the carbonyl oxygen but the difference in base strength is less, and this probably explains why the fractionation factor is reduced to 0.7. To summarise using one kinetic parameter, pk,, the model can explain: ( I ) the size of the rate constants, (2) the slopes of the Brarnsted plots, (3) the range of catalyst pK over which the third-order term is found, (4) the solvent isotope effects and ( 5 ) the carbon hydrogen isotope effects. THE CONCERTED MECHANISM Having established the model, we now enquire under what conditions will the third-order term be observed for other systems? The best chance of observing the third-order term is to choose the conjugate acid-base pair with the PKHA which corresponds to a = +.This PKHA is given by where pKTD describes the equilibrium between the product (the enol) and the reactant (21) (the ketone): PKTU = PK; = 1-pKiH- The optimal PKHA is the mean of pKHA, I and pKHA, 11, and from eqn (18) the concerted transition states exist for a range of PKHA of (~K,)~/dpk,. The route through the concerted transition state is in competition with the HA acid-catalysed route and the A- base route. For this particular acid-base pair (a = 8) the rate constants for these two routes are equal and whereW. J. ALBERY 1589 / 1 \ I I I 1 - 4 - 2 0 2 PH - PK,, FIG.9.-Typical variation with pH of the log of the three contributions to the rate from HA, A- and the third-order term. The acid is the optimal acid for observing the third-order term with kHA = kA-, and we have assumed that the sum of the buffer components is constant: [HA]+[A-] = c. The value ofy is given by eqn (24). Thus the parameter Y gives the largest catalytic advantage that the concerted route has for any acid-base pair. In fig. 9 we sketch the variation with pH of the three contributions to the observed rate for a solution containing a constant concentration of buffer species, c. The separation, y , between the maximum in the third-order term and the sum of the other two terms is given by y = Y+logc-1. (24) At best c/mol dm-3 cannot be much greater than 1 ; hence for the third-order term to be significant we require Y = 1.It can be seen from eqn (23) that the concerted route (large values of Y ) is favoured first for systems with large values of pK,. This is not surprising, since from eqn (16) pKA measures the height of the two ‘unstable’ corners (NW and SE) with respect to the two ‘stable’ corners (SW and NE). The larger the value of pKA the more the system will avoid the unstable corners. Secondly, for the same value of pKA the concerted route is favoured the smaller is the activation barrier given by pk,. However, pKA and pk, will probably be roughly related. Indeed it can be shown that = (pKA)2/32(pkL) = kPk:VA -8HA)2* (25) This equation suggests that systems with large barriers are more likely to have significant third-order terms.Now the rate constant for the A--catalysed reaction [eqn pka = 2Qipk;. (9)] is given by Hence finally we can obtain an estimate for the vital parameter Y in terms of the Brarnsted slopes and the rate constant for the A--catalysed rate, where HA is the This relation therefore allows Y to be calculated from accessible kinetic data.1590 MARCUS THEORY APPLIED TO PROTON TRANSFERS For instance, the ketone/enol transformation of dihydroxyacetone phosphate to glyceraldelyde phosphate catalysed by the enzyme triose phosphate isomerase has been shown to be extremely efficient.ls?l9 One suggested reason for this2*Y2l is that there is concerted catalysis with an enzyme base removing the carbon proton and an enzyme acid protonating the carbonyl group.For such reactions, where the acid and base are both on the same catalyst, the maximum catalytic advantage of the concerted mechanism is given simply by Y. Substitution of typical values in eqn (26) shows that this advantage is likely to be only one or two orders of magnitude. I am grateful to Professors J. P. Guthrie and F. Hegarty and to Dr R. A. More O'Ferrall for stimulating and helpful discussions, and to Mr P. J. Colby for his assistance with the computer calculations. W. J. Albery and J. S. Gelles, J . Chem. SOC., Faraday Trans. 1, 1982, 78, 1569. R. A. Marcus, J. Phys. Chem., 1968, 72, 891. A. 0. Cohen and R. A. Marcus, J. Phys. Chem., 1968, 72,4249. W. J. Albery, A. N. Campbell-Crawford awj J. S. Curran, J . Chem. SOC., Perkin Trans. 2, 1972,2206. J. P. Guthrie and P. A. Cullimore, Can. J . Chem., 1979, 57, 240. J. P. Guthrie, Can. J . Chem., 1979, 57, 797. 'I J. P. Guthrie, Can. J. Chem., 1979, 57, 1177. H. J. Campbell and J. T. Edward, Can. J . Chem., 1960, 38, 2109. N. C. Den0 and M. J. Wisotsky, J. Am. Chem. SOC., 1963, 85, 1735. lo E. M. Arnett, R. P. Quirke and J. W. Larsen, J . Am. Chem. SOC., 1970, 92, 3977. l 1 J. Hine, J . Am. Chem. SOC., 1971, 93, 3701. l2 R. A. McClelland and W. F. Reynolds, Can. J. Chem., 1960,38, 2109. l 3 G. Perdoncin and G. Scorrano, J. Am. Chem. SOC., 1977, 99, 6983. l5 A. F. Hegarty and W. P. Jencks, J . Am. Chem. SOC., 1975, W, 7188. l8 F. H. Westheimer, Chem. Rev., 1961, 61, 265. R. P. Bell and 0. M. Lidwell, Proc. R. SOC. London, Ser. A, 1940, 176, 88. M. M. Kreevoy, T. M. Liang and K. C. Chang, J . Am. Chem. SOC., 1977, 99, 5207. W. J. Albery and J. R. Knowles, Biochemistry, 1976, 15, 5631. I9 J. R. Knowles and W. J. Albery, Ace. Chem. Res., 1977, 10, 105. 2o M. R. Webb and J. R. Knowles, Biochem. J., 1974, 141, 589. 21 M. R. Webb and J. R. Knowles, Biochemistry, 1975, 14, 4692. (PAPER 1/1077)

ISSN:0300-9599    

DOI:10.1039/F19827801579

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J. Chem. SOC., Faraday Trans. 1, 1982, 78, 1591-1594 Adsorption Hysteresis near Saturation Pressure for Nitrogen on Two Carbon Blacks BY W. D. MACHIN Department of Chemistry, Memorial University of Newfoundland, St John’s, Newfoundland, Canada A1 B 3x7 Received 13th July, 1981 Interparticle capillary condensation of nitrogen on standard graphitized carbon blacks, Sterling FT-G (2700) and Vulcan 3-G (2700), becomes extensive near saturation pressure at ca. 77 K. Using a simple model to describe the physical structure of these adsorbents, the relative pressures have been calculated for the fusion of adsorbate interparticle pendular rings. These values are in reasonable agreement with relative pressures corresponding to the lower closure points of each hysteresis loop. Adsorption hysteresis is commonly observed with adsorbents which have high specific areas, but less so with adsorbents which have smaller specific areas.It is customary to regard adsorbents which exhibit hysteresis as being porous and those which do not as non-porous. Experimental requirements for traditional volumetric or gravimetric techniques dictate that the specific area for non-porous absorbents be at least a few m2 g-l. Consequently, hysteresis due to capillary condensation within the spaces between the adsorbent particles should be observable, although it may occur very close to saturation pressure. Brunauer and coworkers1 have reported such a condition for several carbon blacks, one of which displayed hysteresis near saturation, whereas similar hysteresis apparently was lacking with the others.As Everett2 has noted, the phenomenon of hysteresis may not occur with all pore shapes. Some pore shapes that can fill and empty reversibly are, for example, closed tapered cylinders and pendular rings between two particles in contact. Since the latter must exist for virtually all practical adsorbents, including those regarded as non-porous, it is evident that capillary condensation should contribute to the overall adsorption isotherm for ‘non-porous’ adsorbents. In this work the adsorption isotherms for nitrogen on two standard carbon blacks are considered for the region where relative pressures approach saturation. Both systems exhibit extensive hysteresis. EXPERIMENTAL The adsorbents, Sterling FT-G (2700) and Vulcan 3-G (2700), were obtained from the National Physical Laboratory and their properties are described el~ewhere.~ Recommended degassing procedures were closely followed.Research-purity nitrogen (minimum 99.9995 %) was used as supplied by the Matheson Company. The adsorption system is shown schematically in fig. 1. Less than one-tenth of the total line volume was subject to ambient temperature fluctuations, while the major portion of the line was maintained at 297.5 0.05 K. High-vacuum metal valves (Hoke Inc.) were used throughout. A sensitive capacitance manometer (MKS Baratron) was used to determine all pressures. Near atmospheric pressure the limiting precision of the manometer was 0.02 mmHg.* * 1 mmHg x 133.322387 Pa. 15911592 ADSORPTION HYSTERESIS The constant-temperature bath surrounding the sample and the liquid-adsorbate reservoir consisted of a 1 dm3 Dewar flask inside a 4dm3 Dewar flask.Both were filled with liquid nitrogen, and the contents of the inner bath were stirred vigorously. Periodic additions of liquid nitrogen maintained the liquid level within 10 mm of the top of the inner flask. The accumulated error in the amount adsorbed at each point was of the order of +0.4%, this becoming greater for desorption points, especially those where the amount adsorbed was small. Variation in saturation pressure over the course of a run (10-12 h) was CQ. & 1 mmHg, reflecting an average fluctuation of kO.01 K in the bath temperature. Consequently, the error in relative pressure ( p / p o ) at pressures near saturation was ca.kO.002. 4 I I I I I i I FIG. 1.-Adsorption system, (schematic) A, freeze-out; B, bulb, (ca. 200 cm3) thermostatted at 297.5 K; C, capacitance manometer; D, adsorbent; E, liquid-adsorbate reservoir; F, to vacuum and gas handling. RESULTS The relevant portions of the isotherms are shown in fig. 2. At lower relative pressures, the isotherms agree with those published previo~sly.~ Beyond the monolayer region, extending to the onset of hysteresis, the adsorption of nitrogen on both adsorbents is represented adequately by V / Vm = 1.72/[ln which is of the same form as the Frenkel-Halsey-Hill (FHH) equation for multilayer ad~orption.~ V and Vm are the liquid volume of adsorbate at p / p o and the monolayer volume, respectively. The numerical constants in eqn (l), and the lower closure points of the hysteresis loops, were determined graphically.Table 1 presents information relevant to each adsorbent; monolayer capacities and adsorbate volumes have been presented as liquid volumes using the density of liquid nitrogen as given by Baly and D ~ n n a n . ~ Porosities have been calculated from adsorbate volumes at saturation and the adsorbent density as stated by the supplier.6 DISCUSSION Electron micrographs indicate that the primary particles of Sterling FT-G (2700) are polyhedral, approximately spherical, in shape.' The observed porosity of the powdered sample is near that value (0.38) characteristic of randomly packed spheres.8 A randomly packed assembly of spheres should therefore serve as an appropriate simple model for the physical structure of Sterling FT-G (2700).This model may be less satisfactory for the case of Vulcan 3-G (2700), since the small size of the primary particles makes it impossible to obtain direct evidence as to their size and shape. ItW. D. MACHIN 1593 I I I 1 1 I 1 1 0.92 0.94 0.96 0.98 relative pressure, p / p o 0.3 0.2 0.1 FIG. 2.-Nitrogen adsorption isotherms at ca. 77 K; A, Vulcan 3-G (2700); 0, Sterling FT-G (2700); desorption points are solid. TABLE I.-SUMMARY OF RESULTS property Sterling FT-G (2700) Vulcan 3-G (2700) monolayer capaci ty/cm3 g-' B.E.T. area/m2 g-l B.E.T. C value adsorbate volumea/cm3 g-l, porosity adsorbent density/g cm-3 average particle radiusb/nm hysteresis closure point, p / p o relative pressure at a pendular ring filling angle of 30° and contact angle 0' at p / p o = 1.000 4.13 x 10-3 11.6 0.290 f 0.003 850 0.385 0.005 2.14 0.942 0.983 126 2.49 x lop2 70.3 0.840 f 0.005 400 0.627 k0.006 2.0 0.887 0.902 21 a All adsorbate volumes are reported as liquid volumes.Calculated from standard surface area and adsorbent density. 52 FAR 11594 ADSORPTION HYSTERESIS is sufficient for present purposes, however, to consider both adsorbents as assemblages of spherical particles. Based on the Kelvin equation, analysis of the hysteresis loop can provide information, expressed as a pore size distribution, about the physical structure of the adsorbent. The Kelvin equation has been stated in the form9 dv/ds = - ry cos d/RT(ln p / p o ) where dv/ds is the volume-to-surface ratio of those pores which can induce capillary condensation at relative pressure p / p o ; V and y are the molar volume and surface tension, respectively, of the liquid adsorbate at temperature T.An Everett2 and otherslO have noted, the Kelvin equation may be inadequate for low and intermediate relative pressures. In the present case, however, since hysteresis closure occurs at much higher relative pressures, in principle the Kelvin equation should be applicable to the entire hysteresis loop. For example, with a tetrahedral or trigonal array of uniformly spherical particles, capillary condensation should begin as isolated pendular rings at the sphere contact points, and adjacent rings would fuse when the filling angle, as defined by Melrose,ll reaches 30'. On desorption, isolated pendular rings will remain after the bulk of the interparticle condensate has been removed.At this point the relative pressure should again correspond to a pendular ring filling angle of 30' and the lower closure point of the hystereses loop is identified with a return to this condition. Melrose and Wallick12 have presented pendular ring calculations on model systems, and application of their results to the present systems yields the results shown in table 1. While agreement between calculated and experimental values for the lower closure points cannot be considered exact, more refined calculations are not warranted with the present data. Experimental error in the determination of relative pressure shows a pronounced effect on the calculated curvature of the liquid-vapour interface at high relative pressures, amounting to +25% at p / p o = 0.990.CONCLUSIONS Interparticle condensation makes a major contribution to the adsorption of nitrogen on Sterling FT-G (2700) and Vulcan 3-G (2700) at ca. 78 K. Adsorption hysteresis is extensive, and the shape and extent of hysteresis appears consistent with the physical properties of the adsorbents and adsorbate. A more critical analysis of the phenomenon requires measurements determined under more stringent conditions of temperature control. J. Skalny, E. E. Bodor and S. Brunauer, J. Colloid Interface Sci., 1971, 37, 476. D. H. Everett, in The Solid-Gas Interface, ed. E. A. Flood (Marcel Dekker, New York, 1967). vol. 2, chap. 36, p. 1078. D. H. Everett, G. D. Parfitt, K. S . W. Sing and R. Wilson, J. Appl. Chem. Biotechnol., 1974, 24, 199. E. C. C. Baly and F. G. Donnan, J. Chem. Soc., 1902, 81, 907. R. Wilson (Division of Chemical Standards, National Physical Laboratory, 1975), personal communication. ' R. Wilson (Division of Chemical Standards, National Physical Laboratory, 1979), personal communication. P. C. Carman, Flow of Gases Through Porous Media (Butterworths, London, 1956), chap. 1, p. 8. S. J. Gregg and K. S. W. Sing, Adsorption, Surface Area, and Porosity (Academic Press, New York, 1968), chap. 3, p. 137. J. C. Melrose, AZChE J., 1966, 12, 986. * L. Alzamora and J. Cortes, J . Colloid Interface Sci., 1976, 56, 347. lo J. C. P. Broekhoff and W. P. van Beek, J. Chem. Soc., Faraday Trans. 1, 1979, 75, 42. l 2 J. C. Melrose and G. C. Wallick, J . Phys. Chem., 1967, 71, 3676. (PAPER 1 / 1 104)

ISSN:0300-9599    

DOI:10.1039/F19827801591

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J. Chem. SOC., Faraday Trans. I, 1982, 78, 1595-1601 Density, Refractive Index, Viscosity and lH Nuclear Magnetic Resonance Measurements of Dimethyl Sulphoxide at 2 O C Intervals in the Range 20-60 O C Structural Implications BY ROY T. M. BICKNELL, DAVID B. DAVIES AND KENNETH G. LAWRENCE* Department of Chemistry, Birkbeck College, Malet Street, London WCl E 7HX Received 20th July, 1981 The density, refractive index, kinematic viscosity and 400 MHz 'H n.m.r. spectra of pure DMSO have been measured very accurately at two degree intervals in the range 20-60°C. The measurements were analysed statistically to detect marked changes in these physical properties that might indicate a liquid structural transition temperature previously suggested to occur between 40 and 60 OC. Over the temperature range studied it was found that the measurements of density and refractive index vaned linearly with temperature, whereas smooth, curved plots were obtained for measurements of kinematic viscosity and 'H n.m.r.chemical shifts with temperature. There is no evidence from these results that the intermolecular- association structure of DMSO breaks down over the temperature range studied. We have previously measured the relative viscosities of solutions of alkali-metal halides in dimethyl sulphoxide (DMSO) to determine B coefficients of the Jones-Dole equation at 25, 35 and 45 OC.' The magnitude of the B coefficients may include a contribution that arises from molecular-association structure of the solvent. It has been stated2 that 'associated molecules (of DMSO) dissociate between 40 and 60 OC as can be seen from the temperature dependence of several physical properties (density, refraction, increase in limiting conductivity, decrease of the Walden product) '.The temperature dependence of conductivity and Walden product measurements may show anomalies that are caused by the presence of ions and not by an intrinsic property of the solvent. So we examined the available literature of density and refractive index measurements of DMSO for temperature-dependence anomalies and we found that five degrees was the smallest temperature interval at which measurements had been reported (see table 2). Since we were interested in locating the structural transition temperature more precisely than 'between 40 and 60 OC', we decided to measure very accurately the density, kinematic viscosity, refractive index and lH n.m.r.chemical shifts of pure DMSO at two degree intervals over the range 20-60 O C . EXPERIMENTAL SOLVENT The method used to purify the solvent has been described previously3 and only liquid that had an electrolytic conductivity of 4 x S m-l or less was used in these measurements. DMSO is a very hygroscopic solvent so after purification it was stored and transferred to vessels in a nitrogen-filled dry box. 1595 52-21596 STUDY OF DMSO I N THE RANGE 20-60°C DENSITY Densities were measured by the oscillating-cell technique using an Anton Par DMA 601 density cell with a DM60 electronic control. The density, d,, of a liquid measured by the oscillation technique at constant temperature is given by d, = A(t; - f;) + d, where t, and t , are the periodic times for the same number of cell oscillations for the liquid and for water, respectively, d, is the density of water and A is an apparatus constant determined from calibration measurements at the same temperature.The DM60 electronic control displays periodic times for a chosen number of oscillations to the nearest microsecond, so that the degree of precision of the measurements should improve by timing as large a number of oscillations as possible. We found that a precision of + 2 x was conveniently obtained for lo4 oscillations. The cell was calibrated with doubly distilled water and with air immediately before measurements on pure DMSO, and at exactly the same temperature.The density of the air was calculated4 from determinations of the barometric pressure and relative humidity made on each day of the measurements. Water densities were taken from the work of The cell was dried with redistilled acetone and air before filling with pure solvent. g REFRACTIVE INDEX Refractive indices (sodium light) were measured with an Abbe refractometer enclosed in a specially constructed nitrogen-filled dry box. The refractometer was calibrated with 1 - bromonaphthalene. Measurements were made on each of three solvent batches at all temperatures. VISCOSITY The kinematic viscosity, v, of a liquid measured with a capillary tube viscometer is given by v = Ct- K / t (2) where t is the flow time and K and C are apparatus constants. The viscosity measurements were made with a suspended-level viscometer fitted with photoheads and coupled via a relay amplifier to an electronic stopclock.Details of this technique have been described previ~usly.~ The viscometer was calibrated by measuring the flow time with pure water at ten degree intervals from 20 to 5OOC. For the suspended level viscometer the constant C of eqn (2) may be considered independent of temperature,6 and the mean value obtained from the calibration measurements was 5.125 (k0.002) x m2 sP2. The kinetic energy correction constant, K, appeared to increase slightly with temperature, but the increase was within the experimental error, so we have assumed that it is constant for the temperature range studied [ K = 5.4( i0.3) x loP6 m2]. Because of the hygroscopic nature of the solvent and the difficulty of removing the last traces of water from a fairly large viscometer (the calibration is upset by baking in an oven), the viscosity measurements were made for half the range of temperatures with one filling of the viscometer.(These measurements took several days to complete.) The remaining temperatures were then covered after it had been cleaned, dried and recharged with a fresh batch of solvent. The whole series of measurements were duplicated and the mean flow times used to calculate the data presented in this paper. The temperature was maintained constant within iO.005 "C for each measurement for the various techniques by means of a thermostatically controlled water bath, and the temperatures were measured with calibrated thermometers.'H N.M.R. SPECTROSCOPY Pure solvent was sealed in a capillary tube under dry conditions and placed in a high quality 5 mm n.m.r. tube containing [2H6]benzene, which was used to lock the field and whose residual proton signal provides a reference standard. Preliminary 60 MHz lH n.m.r. measurements made on a JEOL FX60 spectrometer indicated a temperature dependence of the DMSO signal, but the accuracy was insufficient for reliable measurements of chemical shift changes of ca. 2.5 x loP3 ppm per degree. Hence measurements of the 400 MHz lH n.m.r. spectra ofR. T. M. BICKNELL, D . B. DAVIES A N D K. G . LAWRENCE 1597 DMSO were made on a Bruker WH400 n.m.r. spectrometer at two degree intervals over the range 26-54 O C . Sixteen scans were collected using 32 k data points over a sweep width of 2400 Hz to give a data resolution of 0.15 Hz per point (0.00038 ppm per point). Three consistent readings taken at 5 min intervals (after 30 min waiting time) were used to indicate thermal equilibrium.It was necessary to re-shim the field at each temperature to improve the accuracy of the measurements. Over the whole temperature range studied the chemical shift of benzene moved only 0.45 Hz (3 data points), whereas the chemical shift of DMSO moved 39.30 Hz (262 data points). Proton measurements of DMSO were also made at each temperature but the change in over the temperature range studied was not large enough to give meaningful results considering the 5-10 7; error involved in each measurement. RESULTS AND DISCUSSION If the structure of DMSO changes with temperature, graphs showing the variation of the physical properties with temperature might be expected to show discontinuities or changes of slope as indicated by Schlafer's graph of his refractive index measurements.' So our data for apparently linear graphs were also examined statistically using standard computer routines to determine whether the results could be represented by two intersecting straight lines.In cases where the physical property did not vary linearly with temperature, the results were fitted to various polynomials of increasing power; best fits were assessed by a consideration of the mean square residuals and by scrutinising the algebraic signs of the residuals for non-random patterns. DENSITY A N D REFRACTIVE INDEX Table 1 gives the results for the density and refractive index measurements of DMSO as a function of temperature.The experimental values of both physical properties were found to fit single straight lines represented by the equations at the foot of the table, and obtained by least-squares linear regression fitting. Columns two and four show the experimental data, and columns three and five the residual between the experimental and calculated values. The residuals for both physical properties are randomly distributed about their respective straight lines and give no indication of a discontinuity over the whole temperature range studied. The precision of the measurements as represented by the mean residual was 1 x lop4 for the refractive index measurements and 1.6 x g for the density measurements.We also submitted literature values of the density and refractive index of DMSO at a number of temperatures to the same statistical examination as our own results (see table 2). The most accurate density measurements are those of Heinrich and Surovyg shown in column 4, and our results compare favourably with theirs. Column 5 shows the residual between their observed values and those calculated with the fitted equation given at the foot of the table. Least-squares regression analysis indicated that their results are best represented by an equation linear in temperature, whereas the results of Clever and Snead8 shown in column 2 are better fitted by a quadratic equation. Schlafer and Schaffernicht's refractive index measurements' (column 8) are ca.0.1 % smaller than ours at each temperature and this may result from differences in the experimental techniques (e.g. our measurements were made under anhydrous conditions in a dry box). These authors also presented a graph of their values as a function of temperature that shows a discontinuity; the experimental points are joined by two almost parallel straight lines for the temperature intervals 20-40 "C and 50-80 OC. We have been unable to find a statistical justification for their discontinuous graph, and suggest that their results are adequately represented by a single straight line. The equation for this line is given at the foot of table 2 with the residuals shown in column 9.1598 S T U D Y OF DMSO I N THE RANGE 20-60°C TABLE I.-VARIATION OF DENSITY ( d ) AND REFRACTIVE INDEX ( n ) OF DMSO WITH TEMPERATURE (?) t/"C d / g ~ r n - ~ Ad n An 20 22 24 26 28 30 32 34 36 38 40 42 44 46 48 50 52 54 56 58 60 1.100 401 1.098 350 1.096 370 1.094 378 1.092 429 1.090 423 1.088 420 1.086 397 1.084 407 1.082 404 1.080 430 1.078 421 1.076 425 1.074 400 1.072 427 1.070 447 1.068 434 1.066 438 1.064 422 1.062 427 1.060 398 12 -41 - 24 - 18 31 22 17 -8 0 -6 18 7 9 - 19 6 24 9 10 -8 -5 - 36 1.4793 1.4785 1.4774 1.4768 1.4760 1.475 1 1.4742 1.473 1 1.4723 1.4715 1.4707 1.4697 1.4687 1.4677 1.4670 1.466 1 1.4649 1.4644 1.4636 1.4629 1.4620 0 0 -2 1 2 1 1 -1 0 1 1 0 -1 -2 -1 -1 -4 0 1 2 2 d = 1.120 3666-9.988 71 x 10-4t.n = 1.488 13-4.391 x 10P4t. Ad and An are residuals. TABLE 2.-LITERATURE VALUES FOR DENSITY ( d / g Cm-3) AND REFRACTIVE INDEX (n) OF DMSO Ad Ad Ad An t/OC da x 104 db x 105 dC x 104 nc x 104 20 1.0980 25 30 1.0913 35 40 1.0816 50 1.0721 60 1.0616 70 80 - - - - -3 1.10046 -8 1.1000 - 1.095 63 4 7 1.090 68 3 1.085 69 -2 - -4 1.08081 4 1.0825 -2 1.07095 6 - - 1 1.060 93 -7 1.0637 - 1.0472 - - 1 1.4783 1.4742 - 3 1.4695 1.4647 .8 1.4600 1.4557 4 1.4517 -1 3 1 -2 -4 -2 4 a Clever and Snead, ref.(8): d = 1.1 103 -4.97 x (-5.3 x t2. Heinrich and Surovy, ref. (9): d = 1.120 30-9.88 x t. Schlafer and Schaffernicht, ref. (7): d = 1.1177 - 8.9 x t ; n = 1.4874 - 4.5 x loP4 t . Ad and An are residuals. KINEMATIC VISCOSITY Measurements of the kinematic viscosity with temperature are summarised in table 3. The data are best fitted by a quartic equation in temperature given at the bottom of the table, and when the results are plotted graphically a smooth curve is obtained that shows no discontinuity.A comparison with the dynamic viscosity values ofR. T. M. BICKNELL, D. B. DAVIES AND K. G. LAWRENCE 1599 TABLE 3.-vARIATION OF KINEMATIC (V) AND DYNAMIC (q) VISCOSITY OF DMSO WITH TEMPERATURE V/CP rllcPU t/OC v/cs residuals x lo4 20 22 24 26 28 30 32 34 36 38 40 42 44 46 48 50 52 54 56 58 60 2-0137 1.9328 1.8583 1.7883 1.7225 1.6599 1.6014 1.5467 1.4948 1.4460 1.4000 1.3562 1.3148 1.2755 1.2383 1.2032 1.1690 1.1370 1.1061 1.077 1 1.0492 3 -6 -1 2 5 -1 -2 0 -1 0 1 0 -1 -1 0 3 - 1 1 -2 1 0 2.2159 - 2.473 - - 1.8100 - - - 1.5126 - _. 1.2880 - - 1.372 - - - 1.1126 - 1.124 1 CS = m2 ssl. 1 CP = kg m-’ s-’.v=3.17344-7.8501 x 10-2t+1.2306x10-3t2-l.116x10-5t3+4.37x10-8t4. l/v = 0.305 41 +8.8026 x lop3 t+4.00 x lo-’ t 2 - 1.1 x lo-’ t3. Schlafer and Schaffernicht, ref. (7). Schlafer and Schaffernicht (column 5) can be made by multiplying our kinematic viscosity and density values at each temperature, and these values are presented in column 4. A number of studies of the influence of temperature on the viscosity of pure liquids have been made. Experimental results and theoretical considerations have led to the (3) equation where q is the dynamic viscosity, E is the energy of viscous flow and classically A is a constant, although EyringlO suggested that it should be a function of temperature. The results for many non- or weakly-associated liquids fit eqn (3), whereas strongly associated solvents like water show a curved plot, as was also found for DMSO.The exponential equation best fitted by our results was of the form In q = In A+E/RT q = a exp (B/T+C/T2+D/T3). (4) On fitting the results to Andrade’P extension of eqn ( 3 ) : In q+i In u = In A’+c/uT ( 5 ) where u is the specific volume of the liquid, a graph less curved than that for eqn (3) was obtained. The computed value of c was 1182 (& 7) cm3 K g-l, which is similar in magnitude to the values of c for other associated liquids investigated by Andrade.ll We also tried a number of alternative empirical equations involving the kinematic1600 STUDY OF DMSO IN THE RANGE 20-60°C or dynamic viscosity and temperature, and their reciprocals. The equation of smallest degree (cubic in temperature) and minimum mean square residual that fitted our results is given at the foot of table 3.CHEMICAL SHIFTS The lH n.m.r. chemical shifts of DMSO at different temperqtures are shown in table 4. Linear regression analysis of these results indicated an approximately linear dependence of chemical shift with temperature with a slope of 1.016 (kO.012) Hz per degree and intercept of 1601.29 (k 0.48) Hz, where the standard errors are little greater than three times the data resolution of the measurement (0.15 Hz per point). The pattern of residuals for linear analysis of the results shown in table 4 (column 3) TABLE 4.-vARIATION OF 'H N.M.R. CHEMICAL SHIFTS OF DMSO WITH TEMPERATUREa chemical shiftb linear regressionC quadratic regressiond tempPC AS/Hz residual/Hz residual/Hz 26 28 30 32 34 36 38 40 42 44 46 48 50 52 54 1627.20 1629.18 1631.62 1634.07 1636.20 1638.19 1640.17 1642.1 5 1644.40 1646.44 1648.08 1649.62 1651.45 1654.05 1656.04 - 0.504 - 0.555 - 0.147 0.272 0.371 0.329 0.278 0.226 0.445 0.453 0.06 1 - 0.430 - 0.632 - 0.063 - 0.104 -0.016 - 0.276 - 0.044 0.229 0.21 5 0.093 - 0.006 - 0.074 0.161 0.2 17 - 0.094 - 0.473 -0.530 0.2 15 0.383 a 400 MHz lH n.m.r.measurements using [2H]benzene as external lock. Chemical shift different in Hz between DMSO and benzene signals. AS = 1.015 75tf 1601.29. AS = 1595.17+ 1.337 t-0.004 t 2 . indicates a slight curvature in behaviour. Non-linear analysis, the results in table 4, indicated a quadratic dependence of chemical shift on temperature as shown by the more random distribution of residuals in column 4.Within the accuracy of the experiment the lH n.m.r. results of DMSO show a smooth variation of chemical shift with temperature over the range26-54 O C and confirm the results ofother measurements in this work that there is no discontinuity in behaviour of DMSO over this temperature range. We therefore conclude that the intermolecular-association structure of DMSO does not change abruptly with temperature in the interval 20-60 OC as previously asserted, and that no structural transition temperature could be identified. We thank Mr Colin Chalmers of the Department of Statistics for help with the computing and statistics, Mr Malcolm Buckingham of the ULIRS 400 MHz n.m.r. Service (Queen Mary College) for making the lH n.m.r. measurements, and the U.L. Central Research Fund for a grant towards the cost of the Anton Paar density apparatus.R. T. M. BICKNELL, D. B. D A V I E S A N D K. G . LAWRENCE 1601 R. T. M. Bicknell, K. G. Lawrence and D. Feakins, J. Chem. SOC., Faraday Trans. I , 1980, 76, 637. * D. Martin and H. G. Hauthal, Dimethyl Sulphoxide, English translation by E. S. Halberstadt (Van Nostrand Rheinhold, England, 1975), p. 52. R. T. M. Bicknell, K. G. Lawrence, M.-A. Seeley, D. Feakins and L. Werblan, J. Chem. SOC., Faraday Trans. I , 1976, 72, 307. Physical Methods of Organic Chemistry, ed. A. Weissberger (Interscience, New York, 3rd edn, 1965), vol. 1, part 1, p. 159. G. S. Kell, J. Chem. Eng. Data, 1975, 20, 97. British Standards Institution Publication BS 188 (B.S.I., London, 1977), p. 12. H. L. Schlafer and W. Schaffernicht, Angew. Chem., 1960, 72, 618. J. Heinrich and J. Surovy, Sb. Pr. Chem. Fak. SVST, 1966, 207. 1941), p. 480. * H. L. Clever and C. C. Snead, J. Phys. Chem., 1963, 67, 918. lo S. Glasstone, K. J. Laidler and H. Eyring, The Theory of Rate Processes (McGraw-Hill, New York, l1 E. N. da C. Andrade, Philos. Mag., 1934, 17, 698. (PAPER 1/1149)

ISSN:0300-9599    

DOI:10.1039/F19827801595

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J . Chem. SOC., Faraday Trans. 1, 1982, 78, 1603-1613 Initial Sintering of Magnesium Oxide in Carbon Dioxide BY TOMOYASU ITO Department of Chemistry, Faculty of Science, Tokyo Metropolitan University, Setagaya, Tokyo, Japan 158 Received 24th July, 198 1 Studies have been made of the initial sintering (surface area diminution and crystallite growth) of MgO freely dispersed in CO,. The sintering rate at 1 123 K was directly proportional to the CO, pressure in the range 0.67-93.1 kPa, and the apparent activation energy for sintering was 153 k 11 kJ mol-'. Increases in the cation vacancy concentration by doping with Mn ions resulted in virtually no variation in the sintering rate. From sintering measurements in 180-enriched CO,, the bulk l80 distribution in crystallites grown by sintering was determined and it was shown that most of the anions migrate only on the surface during sintering.Relative sintering rate constants at 1123 K were (3.4 f 0.3) x lo6, 38 f 8, 1.3 k 0.1 and 1 in H,O, CO,, 0, and Ar, respectively. It is concluded, on the basis of these facts, that sintering is enhanced by increased surface migration of 02- ions (the slower moving ions), caused by repetition of the adsorption-desorption cycle of CO, molecules (anion-exchange mechanism). Imperfections on the surface of polycrystalline oxide powders are greatly affected by the calcination conditions, and only oxide powders prepared under strictly controlled conditions show reproducible surface properties. Previously we have reported'?, the remarkable positive effect of H,O vapour on the initial sintering of freely dispersed MgO powders.This enhanced sintering has been explained by the anion-exchange mechanism, i.e. by increased surface migration of 02- ions caused by repetition of the adsorption-desorption cycle of H,O molecules. A similar effect may be expected for other oxygen-containing molecules. Carbon dioxide is advantageous for such a study because of its high thermal stability and known adsorption states. The main adsorbed species at 773 K has been proved by i.r. spectroscopy to be a bidentate carbonate i ~ n . ~ - ~ Although there have been no reports on the isotopic exchange reaction of oxygen atoms (i.e. the anion-exchange reaction) between CO, and MgO at high temperatures, this reaction proceeds even at 373 K.6 Tomizawa et al.7 have found some accelerating effects of CO, molecules on the initial sintering of MgO powders.However, the atmosphere used by them cannot be considered to be pure CO, because the effects of H,O vapour, evolved during preparation, were ignored. In addition, they proposed no sintering mechanism. In the present paper the initial sintering in CO, is studied and the mechanism is discussed with reference to: (1) the C0,-pressure dependence of the sintering rate, (2) the effect of Mn doping and (3) the I8O distribution in a MgO crystallite. The sintering rates were also measured in H,O, 0, and Ar. EXPERIMENTAL The apparatus and procedure for sintering and adsorption studies have been described previously.2 Carbon dioxide (> 99.99%, Takachiho Chemicals) was predried through a trap kept at 195 K.An isotopic exchange reaction of '*O between 180-enriched CO, and MgO was carried out in the same silica vessel as that used for the sintering study. The concentration of 16031604 l80 in CO, was measured by a mass spectrometer at prescribed time intervals. The initial '*O concentration in CO, was 4.48 atom %. Magnesium oxide specimens were obtained by the thermal decomposition of two varieties of magnesium oxalate dihydrate (MO-6 and M0-9)., Unless otherwise stated, MO-6 was used. Magnesium oxide specimens prepared from MO-9 were usually preoxidized in 13 kPa 0, at 1123 K for 1 h before the sintering run in order to remove any organic contaminants. No essential change in sintering behaviour was produced by the oxygen pretreatment.Magnesium oxalate dihydrates doped with various concentrations of manganese ions were prepared by reaction of magnesium sulphate containing manganese ions with ammonium oxalate. Manganese concentrations in the oxalates, as determined by an atomic absorption spectrometer, were 8, 205, 760 and 1400 atom ppm. I N I T I A L SINTERING OF MgO I N CO, RESULTS INITIAL SINTERING I N PURE MgO The specific surface area, S, of MgO was 268 m2 g-l before the admission of CO,. When MgO was sintered in 0.67-93.1 kPa of CO, at 1123 K, Sdecreased with sintering time, t, as shown in fig. 1. The initial sintering was evidently enhanced by the presence of CO,, but much less than with H,O vapour., The degree of secondary agglomeration between crystallites, S/SX,, where S, is an ideal specific surface area calculated from a crystallite size, was 0.85 k 0.1 throughout the sintering process.Little change in S / S , during the sintering indicates that the surface area reduction was a direct consequence of the crystallite growth. In such a case a kinetic equation S m - S r = k,t (1) is apparently applicable,, where So is S at t = 0, k, is a rate constant for surface area diminution and rn is a constant determined experimentally. Although there is some scatter, the curves in fig. 1 were best fitted to eqn (1) for rn = - 6.0 f 1 .O. The values of k , are shown in fig. 2 as a function of the CO, pressure, pco,; net values obtained by subtracting k, under evacuation (< Pa) from those in the presence of CO, are used. A resultant straight line with a slope of 1.06f0.07 means that k, is directly proportional to pco,.The distribution of micropores in MgO specimens sintered in CO, was obtained by measuring an adsorption isotherm of 0, at 77 K. All the samples examined had no micropores below 1.2 nm diameter and showed the maximum distribution at 2.0-2.5 nm. This suggests that the sintering proceeded in all micropores, and not simply those of a particular diameter. An Arrhenius plot of k , in the range of sintering temperature, T, of 973-1 123 K is shown in fig. 3. The slope gave an apparent activation energy of 153 k 1 1 kJ mol-l for the sintering rate, which seems rather larger than the value of 131 k 18 kJ mol-1 in H,O vapour., I N I T I A L SINTERING OF Mn-DOPED MgO In order to see whether the migration of cation vacancies, VM, is related to the rate-determining step of the sintering process, a sintering study was also carried out on Mn-doped MgO.These specimens were preoxidized in 13 kPa 0, at 1123 K for 1 h before sintering in order to bring most manganese ions to the highest valence state. Values of k, in 20 kPa CO, at 1123 K are shown in fig. 4 as a function of Mn concentration, [Mn], and a least-squares fit gives a slope with 0.06f0.06. In spite of the large variation in [Mn] of 8-1400 atom ppm, k , is substantially constant. E.s.r. spectra of the specimens after sintering showed weak Mn2+ signals. However,T. I T 0 1605 t L I 0 5 10 t l h FIG. 1. 0 10 O 10' lo2 Pco, IkPa FIG. 2. FIG. 1 .-Decrease in surface area of MgO at 1 123 K in CO,.CO, pressure: (a) < (in vacuo), (b) 0.67, (c) 1.33, ( d ) 2.00, (e) 6.65, (f) 33.3, ( g ) 46.6 and (h) 93.1 kPa. FIG. 2.-Plot of log k, against log pco, for the sintering of MgO at 1123 K. Net values obtained by subtracting k , under evacuation from those under the presence of CO,. -i 1O-'L -c h - I 0.0 N E v . -ti 10-15 0.90 0.95 1.00 1.05 1 O3 KIT FIG. 3.-Plot of log k, against T-' for the sintering of MgO in 33.3 kPa CO,.1606 I N I T I A L SINTERING OF MgO I N CO, the quantitative estimation of manganese ions (Mn2+, Mn3+ and Mn4+) by e.s.r. spectra is difficult for highly dispersed MgO.s Therefore, we examined the valence states of manganese ions by measuring the volumetric adsorption of H, at the sintering temperature. The amounts adsorbed, uH, are shown in fig.5 as a function of [Mn], where uH is expressed as the ratio (atom ppm) of the number of hydrogen atoms FIG. 4.-Plot of log k , 0 0 u 0 I 10’ lo2 lo3 1 oL [Mnl (atom ppm) against log [Mn] for the sintering of Mn-doped MgO in 20 kPa CO, at 1123 K. I I I I 0 LOO 800 [ M n ] (atom ppm) FIG. 5.-Plot of uH against [Mn] for Mn-doped MgO at 1123 K. The amount adsorbed is expressed as the ratio of the number of hydrogen atoms adsorbed to that of total Mg*+ ions: (a) after preoxidation (Le. before sintering), (b) after sintering and (c) after the first adsorption of H,. adsorbed to that of total Mg ions. For all the cases uH is roughly expressed as uH = a[Mn]+p. (2) For the specimens after preoxidation [i.e. before sintering, fig. 5(a)], a = 1.8 f 0.6 and p = 890k 300, and for the specimens after the sintering run in CO, [fig.5(b)], a = 1.8 k 0.2 and p = 170 f 70. On the other hand, for the specimens only evacuated after the first H, adsorption [fig. 5(c)], uH is little dependent on [Mn], suggesting thatT. I T 0 1607 most of the manganese ions were in the bivalent state. Therefore, the slopes close to 2 in fig. 5 ( a ) and (b) indicate that most of Mn ions were Mn4+ and that YM originating from Mn4+ ions was present during the sintering in a concentration nearly equal to [Mn]. The difference between fig. 5(a) and (b) may be due to some extrinsic adsorption other than manganese ions, and fig. 5(c) is due to the intrinsic H, adsorption.2 INITIAL SINTERING I N 1 8 0 - ~ ~ ~ ~ ~ ~ CO, If the anion-exchange mechanism is operating it is of interest to measure the rate of isotopic exchange of l80 between 180-enriched CO, and the surface oxide ions.Unfortunately, the exchange rate at the sintering temperature was too rapid, as described below, to observe the rate. On the other hand, information on the bulk l80 distribution of a crystallite sintered in 180-enriched CO, is also useful in discussing E Q 1 - 8 1 I I I 1 0 2 4 1/11 FIG. 6.-Plots of a and D against t for the sintering of MgO in 13.3 kPa 180-enriched CO, at 1123 K. Initial '*O concentrations (atom %) are 4.48 and 0.20 for CO, and MgO, respectively. ths sintering mechanism. For this purpose 13.3 kPa CO, gas containing 4.48 atom T/o l80 was admitted at 1123 K into the sintering vessel with a volume of 26.5 cm3 in which 27 mg MgO containing 0.20 atom % l80 (So = 220 m2 g-l) was present.In fig. 6 the l80 concentration, a, in CO, is plotted as a function o f t . In this figure the crystallite size, D, determined by an X-ray diffraction method,, is also shown. The l80 concentration decreased instantaneously from 4.48 to 2.90 atom % upon the admission of gas, indicating that oxide ions within 1.5 surface layers (ca. 0.3 nm) of the MgO crystallites are very rapidly exchangeable with gaseous CO,. After this, a decreased at a measurable rate. A final equilibrium concentration is calculated to be 1.49 atom % 1 8 0 . If no MgO specimens were present in the vessel, a at t = 4 h was 4.45 atom % l80, which approximately equals the initial value.1608 INITIAL SINTERING OF MgO IN co, 250- 0 5 10 f i h FIG.7.-Surface-area decrease of MgO (origin MO-9) sintered at 1123 K in various gases: (a) in uucuo (< Pa), (b) 20 kPa Ar, (c) 20 kPa 0,, ( d ) 20 kPa CO, and (e) 13.3 Pa H,O. TABLE 1 .-RATE CONSTANTS FOR INITIAL SINTERING OF MgO POWDERS (ORIGIN MO-9) IN 20 kPa ATMOSPHERES AT 1123 Ka gas k,/(m2 g-1)-6 h-l relative rate constant Ar (5.3 k0.5) x 1 0 2 (6.8 k0.5) x 1.3k0.1 co2 (2.0 k 0.4) x 38f8 H20b (1.8k0.2) x (3.4k0.3) x lo6 a Net values obtained by subtracting k, under evacuation from those under the presence of the gases; a value extrapolated from k, in 13.3 Pa H20 vapour. INITIAL SINTERING IN OTHER GASES Information on the initial sintering in atmospheres of oxygen-containing molecules other than CO, may be useful for an understanding of the sintering mechanism.In 20 kPa (13.3 Pa for H,O) of H,O, CO,, 0, and Ar (as a reference), MgO specimens (origin MO-9) were sintered at 1123 K as shown in fig. 7. Rate constants calculated from eqn (1) (rn = - 6.0 f 1 .O) are given in table 1 , where k, for H,O vapour refers to that in 20 kPa, assuming k, to be directly proportional to the pressure of H,O vapour.2T. I T 0 1609 All the rate constants in table 1 are net values obtained by subtracting k, under evacuation from values in the presence of the gases. Relative efficiencies for sintering are in the following order: H,O % CO, > 0, z Ar. DISCUSSION MECHANISM OF INITIAL SINTERING The enhanced sintering is closely related to the adsorption of CO,. We have no reported data on the adsorption at the temperature used in the sintering measurements, but at 773 K volumetric and i.r.spectroscopic investigations have been reported by Evans et al.,3 Gregg et al.4 and Fukuda et al.5 Species adsorbed on MgO at 773 K have been proved by these authors to consist of bidentate carbonate ions /o\ / \ Mg\ 0 P = O as a major component and simple carbonate ions LO 0-cc’ \O as a minor component. Gregg et al.4 also found that at 773 K in 2 kPa CO, an adsorption equilibrium was reached within a few minutes with a surface coverage of ca. 0.06. These results suggest that, under the present sintering conditions, simple and bidentate carbonate species were also formed only in small amounts in a dynamic adsorption equilibrium with gaseous molecules. Possible sintering mechanisms are : (1) a diffusion mechanism (surface, grain- boundary or bulk) ; (2) a viscous flow mechanism ; (3) an evaporation-condensation mechanism ; (4) an adsorption-induced space-charge mechanism and (5) an adsorption- desorption cycle mechanism (anion-exchange mechanism). Of these five, mechanisms (3) and (4) can be excluded., Mechanisms (1) and (2) both depend on the equilibrium adsorption amount, v , of CO, [in mechanism (1) the rate is proportional to v and in (2) to u2], and mechanism (5) depends on the frequency of the adsorption-desorption cycle of CO, molecules under a dynamic adsorption equilibrium.We first look at mechanisms (1) and (2). The adsorption equilibrium of CO, on the surface sites, S’, is expressed as K , S’+CO,eA (3) where Kl is the equilibrium constant and A is the adsorbed species (non-dissociated).(4) From eqn (3), we obtain v = [A1 = K,[S’lP,o2. On the other hand, v can also be expressed as where 8 is the surface coverage and where subscript 0 refers to a state before d In k, adsorption. Then we define A by From eqn (4)-(6) we can obtain A = d In Pco,’ A = 1-81610 I N I T I A L SINTERING OF MgO I N CO, for mechanism (l), since k, cc v. This equation implies that k, cc pe&?). Similarly, for mechanism (2) 3, = 2-26 i.e. k , c c p p ~ : ~ ) . rate-determining, a reaction scheme Next we consider mechanism (5). If, for example, an adsorption process is K . \ S / + C 0 2 2 1 [ 1 k, I + A (9) is applied according to the absolute reaction rate theory, where I is an activated complex (non-dissociated) and K, and k , are the equilibrium constant and the rate constant, respectively.Then the frequency, R, of the adsorption-desorption cycle is and we obtain 3,= 1-0 (1 1) for mechanism (9, since k , cc R. Eqn (1 1) also holds for the case of a desorption process being rate-determining. The experimental values were 8 -g l 4 and 2 = 1.06 0.07 (fig. 2). Therefore, eqn (8) is not applicable to the present case, but either eqn (7) [mechanism (l)] or eqn (1 1) [mechanism (5)] is valid. For ionic oxides, sintering necessitates the migration of both cations and anions, and the sintering rate is determined by the migration rate of the slower of the two. The migration of Mg2+ ions on the surface and in the MgO bulk is not expected to be enhanced by the presence of CO, because non-dissociated adsorbed species (simple and bidentate carbonate ions) lead to no changes in the cation vacancy concentration or in the space charge distribution.This is supported by the sintering on Mn-doped MgO shown in fig. 4 and 5 , where k, is virtually constant against the wide range variation in [Mn], i.e. in VM. Thus cation migration (surface, grain-boundary or bulk) is not considered to be rate-determining. The enhanced sintering in CO, can therefore be attributed to the enhanced migration of the anions (the slower moving ions). Mechanism (1) is unlikely to enhance anion migration since the carbonate ions formed are much larger and much heavier than the original 0,-. Therefore the remaining possible mechanism is (9, the adsorption-desorption cycle.In fact, the isotopic-exchange reaction of lSO between gaseous CO, and surface 0,- by this mechanism proceeds very rapidly (cf. fig. 6, the decrease in a from 4.48 to 2.90 atom % l80). This reaction scheme is represented, for example for a bidentate species, by adsorption /’\ desorption C=O - Mg2+ +02-+ COO, (12) Mg\o/ Mg2+ + 0;- + CO, - where the subscript c refers to the 0 atom originally belonging to the oxide. The anion formed by the desorption of the COO, molecule will be in the site next to the original 0;-. This results in migration of the anion on the surface. Guilliatt et a1.8 found that their Mn-doped MgO was sintered much faster in 0, than in H, or in vacuo, and they concluded that cation diffusion by the vacancy mechanism was rate-determining. In the present study, however, surface migration of anions has been proved to be rate-determining.This apparent difference is believed to result mainly from the difference in sintering temperature. Sintering in flowing 0,T. I T 0 161 1 proceeded rapidly only above 1 173 K,8 while in the present study sintering was carried out at 1 123 K. In general, the rate of migration of the cations exceeds that of the anions in one temperature range but the reverse is true in a different range, since the two ions may migrate by different mechanisms. DISTRIBUTION OF l80 IN MgO When the size of a crystallite increases from Do to D by sintering, each bulk anion in this region exists as a surface ion of the growing crystallite for a certain period during the sintering. In 180-enriched CO, the concentration of l80 in the surface lattice anions is equal at any time to that in gaseous CO, because the exchange reaction between the two proceeds instantaneously.Therefore, the relation between a and D in fig. 6 indicates the distribution of 1802- ions in the bulk of the MgO crystallite grown by 3.0 1 0 1 2 3 4 5 r/nm FIG. 8.-Plot of a against r for a crystallite of MgO sintered in 13.3 kPa I80-enriched CO, at 1123 K. Initial l80 concentrations (atom x) are 4.48 and 0.20 for CO, and MgO, respectively. Upon the admission of CO, the anions within 1.5 surface layers (3.05 < r/nm < 3.35) are instantaneously exchangeable. TABLE 2.-NUMBER OF 1802- IONS IN MgOa 0 0.5 1 .o 2.0 4.0 r/nm 3.35 3.77 3.98 4.30 4.61 Sx/m2 g-l 245 21 8 206 191 178 alatom % (4.48) 2.54 2.41 2.28 2.13 n/ 1 016b 2.46 1.73 1.47 1.16 0.94 q// 1 0lBC (0.80) 4.98 5.5 1 6.00 6.28 T, 1 1 0 led - 4.18 4.40 4.61 4.85 GlTd - 1.19 1.25 1.30 1.29 a With estimated errors of r ( f 0.08 nm), S, ( & 4 m2 g-l), a ( f 0.02 atom %), n (k 0.03 x 10l6) and T, and Td (f0.15 x lo1*); the number of crystallites: the number of 1802- ions in all crystallites evaluated from the distribution curve in fig.8 (it is assumed that 15% of surface anions are not exchangeable since the seondary agglomerations between crystallites are 0.85 k 0.1); the number of 1 8 0 2 - ions in all crystallites evaluated from the reduction in l80 concentration in gaseous CO,.1612 I N I T I A L SINTERING OF MgO I N CO, sintering if the bulk diffusion of anions is negligible. Here we assume cubic crystallites with a uniform size’ and define Y (= D / 2 ) as the distance from the centre of the cube to the surface plane.A plot of a against Y in a crystallite is shown in fig. 8, which assumes that anions within 1.5 surface layers are very easily exchangeable upon admission of CO,. If the l 8 0 distribution curve in fig. 8 reflects the real situation, the number of 1 8 0 2 - ions in all crystallites (equal to the number of 1 8 0 2 - ions in a crystallite mutiplied by the number of crystallites) evaluated from the distribution curve must be nearly equal at any time to that calculated from the reduction in l80 concentration in CO,. Calculated results are shown in table 2, which reveals that the ratios, ZJTL, of the values obtained by the two methods are 1.2-1.3 irrespective of t.This constancy and insignificant deviation from 1 suggest the approximate validity of the distribution curve in fig. 8. (If anions diffused freely in bulk, T,/ Ti would become 2.4 and 1.7 at t = 0.5 and 4.0 h, respectively, showing a non-constancy and larger deviations from 1 .) These facts indicate that the assumption of insignificant diffusivity of the bulk anions is valid and that most anions migrate only in a surface layer, which provides support for the anion-exchange mechanism in the initial sintering. SINTERING RATE IN VARIOUS GASES As shown in table I the relative sintering rate constants are (3.4 0.3) x lo6, 38 f. 8 and 1.3 0. I for H,O, CO, and 0,, respectively. Since the sintering in H,O vapour also proceeds by an anion-exchange mechanism,, the isotopic exchange reaction rate of the surface oxide ions in H,O vapour should be ca.lo5 times that in CO,. Unfortunately the exchange rates were too rapid to determine the rate constants.2 In addition there are no reported studies of this subject except for the case of oxygen molecules around 750 K 9 7 lo The theoretical calculation of the exchange rate constant is fairly difficult because the detailed mechanism is obscure. However, the observed pressure dependence, A, of the sintering rate constant is 1 in both H,02 and CO,. This suggests that a molecular adsorption process is rate-determining for the sintering reactions (as for the exchange reactions). In such a case, the exchange rate can be approximately predicted by the theory of absolute rates.ll For CO, an adsorption process is expressed by eqn (9) and the adsorption rate under a dynamic adsorption equilibrium is represented by R=-- kT ‘I exp ( - E / k T ) [CO,] [S’] h qco* 4 s where k is Boltzmann’s constant, h is Plank’s constant, E is the activation energy of adsorption at absolute zero and q is the molecular partition function per unit volume or unit area.For the simplest case qI and qsl may be equated to unity, and [S’] is assumed to be [S’], since 9 6 1. Using as E the apparent activation energy observed for sintering, R can be very roughly estimated. A similar calculation for H,O leads to the result that the adsorption rate for H,O under sintering conditions is ca. lo3 times that for CO,. Although semi-quantitative, this is consistent with the observed sintering rates.In addition, the known properties of these gases indicate that the extents of chemical interactions with MgO surfaces at high temperatures are in the order H,O > CO, > 0,. For example, reported adsorption amounts in terms of apparent surface coverage are as follows: 0.1 512 in 0.8 kPa H,O vapour at 923 K and @.313 in 0.6 kPa H,O vapour at 773 K, 0.064 in 2.0 kPa CO, at 773 K and ca. 0.000214 in 0.008 kPa 0, at 773 K. The same order may be applicable to the isotopic exchange rate. These considerations may give support to the theory that sintering in these atmospheres proceeds by an anion-exc hange mechanism.T. I T 0 161 3 I thank Prof. Taneki Tokuda and his collaborators at the Tokyo Metropolitan University for helpful discussions and suggestions during the course of this work. T. Ito and T. Tokuda, Nippon Kagaku Kaishi, 1974, 248. T. Ito, M. Fujita, M. Watanabe and T. Tokuda, Bull. Chem. SOC. Jpn, 1981, 54, 2412. J. V. Evans and T. L. Whateley, Trans. Faraday Soc., 1967, 63, 2769. S. J. Gregg and J. D. Ramsay, J . Chem. SOC. A, 1970, 2784. Y. Fukuda and K. Tanabe, Bull. Chem. SOC. Jpn, 1973, 46, 1616. 0. V. Krylov, Z. A. Markova, I. I. Tretyakov and E. A. Fokina, Kinet. Katal., 1965, 6, 128. I. F. Guilliatt, and N. H. Brett, J. Chem. SOC., Faraday Trans. I , 1972, 68, 429. E. R. S. Winter, J. Chem. SOC. A, 1968, 2889. A. Clark, The Theory of Adsorption and Catalysis (Academic Press, New York, 1970), p. 209. London, 1957), vol. 2, p. 309. ' T. Tomizawan, H. Hashimoto and K. Moteki, Kogyo Kagaku Zasshi, 1966, 69, 2263. lo J. Novakova, Catal. Rev., 1970, 4, 77. l 2 A. G. Oblad, S. W. Weller and G. A. Mills, Proc. 2nd Znt. Congr. Surface Actiuity (Academic Press, l3 T. Ito, K. Kanehon and T. Tokuda, Z . Phys. Chem. (N.F.), 1976, 103, 203. l4 R. J. Breakspere and L. A. R. Hassan, Aust. J. Chem., 1977, 30, 971. (PAPER 1 / 1 178)

ISSN:0300-9599    

DOI:10.1039/F19827801603

出版商:RSC    

年代:1982

数据来源: RSC