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Oxygen-18 exchange reactions between gaseous oxygen and certain oxygen-containing inorganic salts
J. Inorg.Nucl. Chem., 1965, Vol. 27, pp. 2161 to 2169. PergamonPress Ltd. Printedin Northern Ireland
OXYGEN-18 EXCHANGE REACTIONS BETWEEN GASEOUS OXYGEN AND CERTAIN OXYGENCONTAINING INORGANIC SALTS* B. Z. SHAKHASHIRIand G. GORDON Department of Chemistry, University of Maryland, College Park, Maryland
(Received 27 January 1965) AImtraet--Rate measurements have been carried out on the 180 exchange reaction between oxygen gas and certain alkali metal chromates, molybdates, tungstates, phosphates and sulphates in the 600-850°C temperature range. It was observed that the rate of exchange was influenced by the nature of the cation and the nature of the central atom. In the alkali metal chromate series, the potassium salt exhibited the least tendency to exchange, but the activation energy for the exchange process was observed to decrease with an increase in the cation radius. In contrast, both the rate of exchange and the activation energy decreased as the effective nuclear charge of the central atom increased, as the central atom-oxygen bond distance decreased, and as the oxidation state of the central atom increased. A mechanism is proposed for the exchange reaction and a correlation between the rate of exchange and various thermal and thermodynamic properties is presented. INTRODUCTION
TH~ xsO exchange reactions between gaseous oxygen and solid oxides have been widely investigated by several workers. (x) These investigations have shown in general that exchange can take place if the temperature is above 300° and that the initially rapid exchange which does take place on the surface is followed by a slow exchange, whose rate is controlled by the diffusion of oxygen from the inner layers of the lattice to the surface. SPITSYN and FINIKOVt2) have studied the isotopic exchange between gaseous oxygen and the alkali metal sulphates in the 680-820 ° temperature range. They found that the exchange proceeds not only on the surface but also in the inner layers of the crystals. The activation energies for the exchange processes with lithium, sodium, potassium, rubidium and caesium sulphates were 39, 54, 57, 40 and 24 kcal/mole respectively and were based on the assumption that the limiting step occurred on the surface. It was shown that the activation energy depends on the nature of the cation and that marked variations in the relative strength of the oxygen bonds do exist. This study was undertaken in an attempt to determine the effect of various salt parameters, such as the effective nuclear charge of the central atom, the central atom-oxygen bond distance, and the oxidation state of the central atom on both the rate of exchange and the activation energy. * This work was supported by the United States Atomic Energy Commission under grant No. AT-(40-1)-2858 and is based on a thesis submitted by B. Z. S. to the Graduate School of the University of Maryland in partial fulfillment for the degree of Master of Science (1964). (t) For example, see E. R. S. WINTER, Advances in Catalysis, X, pp. 196-241. Academic, New York
(1958). (2) V. I. SPITSVNand V. G. FINIKOV, Dokl. Akad. Nauk SSSR. 108, 491 (1956). CA 51:3249h. 2161
2162
B . Z . SHAKHASHIRIand G. GORDON EXPERIMENTAL
General procedure The technique employed for studying the isotopic exchange between gaseous oxygen and the various salts consisted of placing a 30-50 mg sample of an ~80-enriched salt (0.60-1.23 ~ xsO) in a Vycor or quartz tube contained in a small furnace. The sample was degassed for a minimum of one hour (<10 -4 torr) at the temperature at which the exchange reaction was to be studied. Approximately 5000 torr-ml of oxygen gas of normal isotopic composition (0.20 ~ ~80) was introduced into the reaction vessel. A small capillary tubing connected the reaction vessel to the inlet system of a Nuclide Analysis Associates RMS-11 Isotope Ratio Mass Spectrometer and allowed a continuous flow of gas. The total volume occupied by the gas was about 100 ml. The progress of the reaction was monitored mass spectrometrically by periodically measuring the change in the ratio of the mass 34-32 (xnO180 to leO160) peaks. Samples of the enriched salts were heated for two hours at various temperatures from 500-850 ° in the exchange vessel which was open to the mass spectrometer. The extent of decomposition, if any, in the temperature range of the study of each salt was checked by comparing the mass spectra of the residual gases found in the vacuum line before and after the heat treatment and by observing any loss in weight of the sample.
Materials All the salts used in this study were enriched in 180. The method of enrichment depended on the nature of the oxy-anion, on the relative ease with which it exchanged oxygen with water, and on the solubility of the resulting compound. In all methods water from the Isomet Corporation which contained 1"55 % xsO was used. Chromium (VI) oxide. Ten grammes of anhydrous chromium (VI) oxide was dissolved in 18 ml of xsO-enriched water in a glass stoppered flask and allowed to stand for 1 hr, since it has been shown ~8~ that isotopic equilibrium is reached very quickly. Portions of the resulting oxygen-18 enriched chromate solution were used directly in the following preparations. Lithium chromate. Two and one half grammes of anhydrous lithium hydroxide was dissolved in 8 ml of distilled water. Excess ~80-enriched chromate solution was added to destroy any lithium carbonate that might have been present. In order to insure the presence of mononuclear chromium species, solid lithium hydroxide was added ¢4~ until the pH of the solution was 8. The solution was evaporated slowly in this and all of the following preparations in order to obtain the appropriate pure salt. Sodium chromate and potassium chromate. BALOGAand EARLEY15~report a half-life of 17 hr for the oxygen exchange reaction between the chromate ion and water at 25 °. Therefore, 17-5 g of anhydrous sodium chromate was dissolved in 20 ml of xsO-enriched water and allowed to stand for 6 days in a glass stoppered flask. Similarly, 11.6 g of anhydrous potassium chromate was dissolved in 20 mi of 1sO-enriched water and allowed to stand for 6 days in a glass stoppered flask. Rubidium chromate and caesium chromate. Twelve grammes of anhydrous rubidium carbonate and 20 g of anhydrous caesium carbonate was treated separately with stoichiometric amounts of the asO-enriched chromate solution. The final pH was 8. Sodium molybdate and sodium tungstate. HALL and ALEXANDEac~ observed that the exchange reactions of oxygen between water and both molybdate and tungstate ions were fast. Ten grammes of anhydrous sodium molybdate was dissolved in 18 ml of 180-enriched water and allowed to stand 10 days in a glass stoppered flask. The pH was adjusted to 8 by means of a carbonate-free sodium hydroxide solution in order to insure the presence of mononuclear species of molybdenum/v~ Similarly, 10 g of anhydrous sodium tungstate was dissolved in 18 ml of 1sO-enriched water and allowed to stand for a minimum of 10 days in a glass stoppered flask. The pH was adjusted to a ~s~ I. DOSTROVSKYand D. SAMUEL,Inorganic Isotopic Syntheses, R. H. HERBER (Editor), p. 139. Benjamin, New York (1962). la) y . I. SASKI, Acta Chem. Scand. 16, 719 (1962). ~ M. R. BALOGAand J. E. EARLEY,J. Phys. Chem. 67, 964 (1963). ~n~N. F. HALL and O. R. ALEXANDER,J. Amer. Chem. Soc. 62, 3455 (1940). tT) y . I. SASKI, I. LINDQVISTand L. G. SILLEN,J. Inorg. NucL Chem. 9, 93 (1959).
Oxygen-18 exchange reactions between gaseous oxygen and certain salts
2163
value of 10 by means of a carbonate-free sodium hydroxide solution in order to insure the presence of mononuclear species of tungsten.C8 Potassium phosphate. Fourteen grammes of anhydrous potassium dihydrogen phosphate was dissolved in 27 ml of 180-enriched water and the solution was heated under reflux at 90 ° for 6 days. The solution was cooled to room temperature and the pH was adjusted to 13 by means of a potassium hydroxide solution. Potassium sulphate. HALL and ALEXANDERt°) and WINTER and co-workers, ~9~ have shown that the sulphate ion and water do not exchange oxygen while the rate of oxygen exchange between sulphite ion and water was measurable. Therefore, 8 g of anhydrous potassium sulphite was dissolved in 14 ml of 180-enriched water and allowed to stand overnight in a glass stoppered flask. The solution was electrolysed by using a sheet of platinum of large surface area (about 6 cm ~) as the anode and a platinum gauze as the cathode for a period of at least 6 hr at a current density of about 0.3 A/cm ~. It was necessary to add a few milliliters of lsO-enricbed water to the cell to compensate for evaporation losses. The amount of sulphite left was determined to be less than 1 ~ by testing with a standard permanganate solution. The remaining sulphite was removed by boiling the solution after the addition of a few drops of concentrated sulphuric acid. Potassium perchlorate. HOERING and co-workers t1°~ have estimated the half-life for the oxygen exchange reaction between the perchlorate ion and water at room temperature to be greater than 100 years. They also observed that the rate of the oxygen exchange reaction between the chlorate ion and water is measurable at 100 °. Eleven grammes of anhydrous sodium chlorate was dissolved in 27 ml of 1sO-enriched water, to which 1.2 ml of concentrated perchloric acid was added. The solution was heated under reflux overnight in order to allow isotopic equilibrium to obtain and then it was electrolysed for a period of 7 hr. Analysis showed 99.9 ~ conversion of chlorate to perchlorate. The resulting solution was treated with concentrated potassium nitrate solution to precipitate the potassium perchlorate. In all cases, the freshly precipitated salt was collected by filtration, washed once with ethyl alcohol or acetone and once with ether. Each salt was dried in an oven for 2 hr at 400 ° and dried further by heating for an additional 2 hr under vacuum at a temperature well below the melting point or decomposition temperature. Oxygen and carbon dioxide of normal isotopic composition (0'20 ~ asO) were obtained from the Matheson Company.
Isotopic analysis of oxygen in the salts The oxygen isotopic analysis of the salts was carried out by converting the oxygen into a gaseous form which was more suitable for mass spectrometry. Since in general the decomposition of the salts takes place at temperatures in excess of 1000 °, the salts were reduced with mercury (II) cyanide to form carbon dioxide with the exception of potassium perchlorate and lithium chromate which were decomposed at 550 ° and 700 ° respectively. The method of reduction to carbon dioxide was that of ANBAR and GUTTMANN.t11) The samples were heated for 2 hr at 400 ° with a mixture of mercury (II) chloride and mercury (II) cyanide; the carbon dioxide formed was purified over zinc amalgam at 200 ° . The isotopic composition of the carbon dioxide was determined mass spectrometrically by measuring the mass 46/(44 + 45) ratio (12ClaOXsO/lzCaeOa60 + 13C160160). The contribution of the x~O isotope is negligible for these samples which contain less than 1-55 ~ 180. The enrichment of the salt was calculated from the measured ratio for the enriched and normal samples.
Apparatus The exchange vessel consisted of a Vycor or a quartz tube 22 m m × 21 cm equipped with a 24/40 standard taper joint. The salt sample was weighed in a small Vycor tube (13 m m × 2 cm) and was placed in the exchange vessel. The vessel was evacuated on a high vacuum line which was equipped with a two stage mercury diffusion pump and a rotary type oil pump. The pressure was monitored ts~ y . I. SASKI, Acta Chem. Scand. 15, 175 (1961). tg~ E. R. S. WINTER, M. CARLTON and H. V. A. BRISCOE, J. Chem. Soc. 131 (1937). c10~T. C. HOERING, F. T. ISH/MORI and H. O. McDONALD, J. Amer. Chem. Soc. 80, 3876 (1958). iltl M. ANBAR and S. GUTrMANN, Int. J. Appl. Rad. Isotopes 5, 233 (1959).
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B.Z. SHAKHASHIRIand G. GORDON
by means of RCA 1946 Thermocouple Gauge. The sample was heated by a small furnace which was constructed in the following manner: Chromel "A" resistance wire was wound about the vessel, covered with asbestos and gypsum, and connected to a Variac; the resistance of the furnace varied from 15 to 29 f~ and the total length of the furnace was about 14 cm. A more elaborate furnace was constructed as follows: chromel "A" resistance wire (80 f~, 150 W at 110 V) was wound about the vessel and covered as before and connected to a Variac through a temperature regulator, two wires of lower resistance were connected in parallel (46 t), 270 W at 110 V) to another Variac. It was found that the second type furnace reached the desired temperature much more quickly than the first type, but the variation in equilibrium temperature of both furnaces was less than 4-0.5 °. The actual temperature control was maintained by means of a Model 283C Barber-Coleman Temperature Regulator. The temperature was measured by means of an insulated chromel-alumel thermocouple obtained from Advanced Technology Laboratories which was placed through a small Vycor tubing which extended into the middle of the exchange vessel. The output of the thermocouple was determined with a Calibration Standards Corporation VA-100A Precision Volt-Amp Meter to a precision of +0.01 mV. RESULTS Vycor and quartz were shown not to exchange oxygen with gaseous oxygen at 950 °. F o r example, after a period o f two hours, the change in the enrichment o f an 8 fold enriched (12) sample o f gaseous oxygen was less than 0.3 ~o. This is in agreement with HUTCHINSON (13) who reports no detectable exchange in the silica-oxygen system below 1000 °. The results o f individual exchange experiments were analysed by means o f the McKAY equation, c14a5) In (1 -- Ft) = --[(a + b)/(ab)](Rt)
(1)
where a and b are the n u m b e r o f moles o f exchangeable oxygen contained in the salt and gas respectively, R is the rate o f exchange in mole sec -1, t is the time in seconds and F t is defined as the fraction F , = ( N t - - No)/(N~o - - N t )
(2)
at any time t. The quantity N t is the measured 180 to 160 ratio o f the gas at any time t, N Ois the initial ratio, and No~ is the ratio at isotopic equilibrium. No~ was calculated according to Equation (3) Noo =
n,(No)i
ni
(3)
0
where i is the n u m b e r of oxygen-containing components in the system and n is the n u m b e r o f moles o f oxygen. The validity o f Equation (3) was checked experimentally. The agreement between the calculated and measured ratio was better than 0.4 ~o (for example, calc. 4.98 fold; measured, 5.00 fold). The calculation o f the rate o f exchange requires an exact knowledge o f the surface area o f the salt and the n u m b e r o f moles o f oxygen in the gas phase. The ideal gas law was used to calculate the moles o f gas in the gas phase; however, the surface areas o f the salts at the temperatures o f the reactions are not known. Therefore, an apparent (12)Oxygen of normal isotopic composition would correspond to an enrichment of 1-00 on this scale. t18) D. A. HUTCHINSON,J. Chem. Phys. 22, 758 (1954). (14) H. A. C. McKAY, Nature, Lond. 142, 997 (1938). (~5)All of the calculations were done on the I.B.M. 7094 Computer in the University of Maryland Computer Science Center.
Oxygen-18 exchange reactions between gaseous oxygen and certain salts
2165
rate constant, k, was defined as [R(a + blab)] and was obtained directly from the slope of the McKAy plot. t14) Similarly, an apparent activation energy, E~, was obtained from the slope of a graph of log k as a function of the reciprocal of the absolute temperature. Although neither k nor E , are absolute quantities, they can be useful for the purpose of comparison of the effect of variation in the salt parameters discussed earlier.
I'0 0'9 0'8 0'7 0"6
0'5 I-F 0'4
0,3
0"2
I 600
I 1200
I I 1800 2400 T I M E (SECONDS)
I 3000
I 3600
I 4,¢00
FIG. 1.--Graph of log (1 -- F) as a function of time for the K~CrO~(s)~O2(g) exchange reaction at 855°. The results of a typical experiment are shown in Fig. 1. Presumably, the fast rate of exchange is that taking place at the surface of the salt while the slower rate is controlled by diffusion from the inner layers of the salt to the surface. This is in agreement with the studies of WINTERtl~ on the reaction between gaseous oxygen and metal oxides and the studies of SPITSYN and FINIKOVt~) who report that the exchange process between gaseous oxygen and the alkali metal sulphates proceeds not only on the surface, but also takes place in the inner layers of the salt. In this study the extent of the exchange reaction was followed for values of F < 0.35 which corresponds to the linear region of the McKAY plot. The individual values of k were reproducible to better than 3 per cent and the standard deviation in the apparent activation energy was no greater than q-3 kcal/mole, except in the case of sodium chromate where the deviation was + 10 kcal/mole. The apparent rate constants and apparent activation energies for the oxygen exchange reaction between gaseous oxygen and the alkali metal chromates are shown in Table 1. These data were used to calculate apparent rate constants at other temperatures and those at 623 ° are also recorded in Table 1.
2166
I.
Z.
TABLE 1 . - - T H E EFFECT
SHArd~ASHI~ and G. GORDON OF THE CATIONON THEHETEROGENEOUSOXYGEN EXCHANGE REACTION
Salt
(sec-1)
e2a ° X 105 kenle" (sec-t)
Li2CrO4 Na2CrO4 K2CrO~ Rb2CrO4 Cs2CrO~
6"64* 6"86t 1"58~ 12"5§ 7"89¶
6"64 0.007 0.004 1"88 1.69
kmeas • >( 1 0 5
• 623°
~ 712°
~ 741 °
Temp. Range E~ (°C) (kcal/mole) 598-623 680-712 723-741 713-745 675-710
§ 713°
152 135 92 37 31
¶ 710°
The alkali metal chromates were observed to change colour f r o m yellow to red when heated above 150 °. The colour change was reversible except in the case o f lithium chromate which developed a deep b r o w n colour u p o n cooling to r o o m temperature. The results o f the oxygen exchange reaction between gaseous oxygen and various other sodium and potassium salts are shown in Table 2. TABLE 2.--THE EFFECT OF THE CENTRAL ATOM ON THE HETEROGENEOUS OXYGEN EXCHANGE REACTION 778 ~ keale ' x 105 (sec-a)
kmeas" X 10 5
Salt Na2CrO4 NazMoO4 Na2WO4 K3PO4 K~SO4 KC10~ * 712°
(sec-1)
6"86* 521 l'50t 1-50 0.57:~ 0"18 5"10 § 21"5 0"52¶ 0"97 decomposed without exchange[I ~ 778°
~ 808°
§ 740°
Temp. Range E~ (°C) (kcal/mole) 680-712 746-810 808-845 727-767 741-851 ¶ 741°
135 94 85 80 36
Jl 550°
DISCUSSION The effect o f the cation The alkali metal chromates and sulphates (except for lithium and sodium) have been observed to be isomorphous (16-1s) and each c h r o m i u m a t o m is tetrahedrally surrounded by four oxygen atoms at an average distance of 1.60 A. I f the apparent rate constants can be used as a measure o f the order o f lability with respect to oxygen exchange, then the order for the alkali metal chromates at 623 ° is lithium > rubidium ~ caesium > sodium > potassium as shown in Table 1. This is in agreement with the observations o f SPITSYN and FINIKOV. t2} At 710 ° the same relative order obtains (neglecting lithium chromate which would decompose at 700°). Even t h o u g h the apparent activation energy m a y change with large variations in temperature, it is interesting to note the apparent rate constants at 740 ° where the lability o f sodium chromate becomes nearly equal to that o f rubidium chromate. j. W. MELLOR, A Comprehensive Treatise on Inorganic and Theoretical Chemistry, Vol. XI, pp. 245, 251. Longmans, Green, New York (1931). (xT~H. W. SMITHand M. Y. COLBY,Z. Krist. 103, 90 (1941). (18~j. j. MILLER,Z. Krist. 99, 32 (1938). lle)
Oxygen-18exchangereactionsbetweengaseousoxygenand certain salts
2167
It is generally accepted that the polarizing power of ions increases with increasing positive charge but decreases with increasing size. Observations that salts of oxyanions with large cations tend to be more stable thermodynamically than salts of the same oxy-anion and small cations suggest that the difference may be associated with the polarization of the anion by the cation, ttg) Therefore, as the degree of polarization increases, the stability of the anion should be expected to decrease. On the other hand, SPITSYN(~) has suggested the possibility of mutual polarization between all of the atoms which form the oxy-anion salt. The relatively strongly polarizing lithium and sodium ions would be slightly polarized by the oxygens of the anions, but the relatively weakly polarizing rubidium and caesium ions would be more polarized by neighbouring oxygens, which would result in a weakening of the central atom-oxygen bond. Since these effects oppose each other, it is indeed possible that potassium salts could show the least tendency to undergo oxygen exchange if the extent of these interactions were maximized. This interpretation would be in agreement with other results tz~ which suggest that the potassium salt does exhibit the least tendency to exchange. The standard heats of formation of the alkali metal sulphates increase with potassium > lithium > rubidium > caesium > sodium, t~l~ This order is markedly different from that observed for the oxygen exchange labilities. The effect of the bond distance The central atom-oxygen bond distance in the isomorphous potassium chromate and sulphate t ~ are 1-60 and 1.49 A respectively. The apparent rate constants shown in Tables 1 and 2 indicate that the chromate anion is more labile with respect to oxygen exchange than is the sulphate anion. This is in agreement with the thermodynamic stability of these salts in that the standard heat of formation of potassium sulphate is larger than that of potassium chromate.t2~ The effective nuclear charges of the central atoms of these anions differ only slightly even though the atomic numbers of sulphur and chromium are 16 and 24 respectively. However, sulphate and chromate differ markedly in their central atom-oxygen bond distances. This suggests that perhaps the apparent rate constant for the heterogeneous oxygen exchange reaction can be used as a measure of the relative strength of the central atom-oxygen bond. It is interesting to note that the rate of oxygen exchange between water and the two anions follows the same trend. ~5'6~ The effect of effective nuclear charge of the central atom The apparent rate constants shown in Table 2 suggest that the order of lability with respect to oxygen exchange is chromate > molybdate > tungstate which is the same as the order of decreasing thermodynamic stability.C21~ The observations of SeITSYN and KULESHOV{~°'~ indicate that sodium tungstate is also more stable thermally than the isomorphous sodium molybdatet24'~5~ in that the loss in weight t19~R. T. SAI'a~RSON,ChemicalPeriodicity, pp. 162-166. Reinhold,New York (1960). t~o~V. I. SPITSX,N, Zhur. Obsch. Khim. 17, 11 (1947); 20, 550 (1950); CA 44:6243c. t21}W. M. LATIMER,Oxidation Potentials (2nd ed.), pp. 329-335. Prentice-Hall,New York (1952). {~ M. A'ro~tand R. E. Rtn,rDLE,J. Chem. Phys. 29, 1306 (1958). t~a~V. I. SerrsvN and I. M. KUL~HOV,Zhur. Obsch. Khim. 21, 401 (1951); CA 45:5553i. ~24}I. LINDQVIST, hcta Chem. Scand. 4, 1066 (1950). 12s~A. F. WELLS, Structural Inorganic Chemistry (3rd ed.), p. 445. Clarendon, Oxford (1962). 3
2168
B.Z. SrIAKHASrmUand G. GORDON
after heating for 6 hr at 1200° was 2.30 ~o for the molybdate and 0.59 ~o for the tungstate. Molybdenum and tungsten, are tetrahedrally surrounded by four oxygen atoms at an average distance of 2.0 A, but the chromium-oxygen bond distance is 1"60 A. Apparently, the increase in the effective nuclear charge of the central atoms has a greater influence than the increase in central atom-oxygen bond distance in that sodium tungstate is more stable thermodynamically, thermally and with respect to oxygen exchange than the corresponding molybdate and chromate. The effect o f the oxidation state o f the central atom
The isoelectronic phosphate, sulphate and perchlorate anions are tetrahedralt22'26,zT~ with central atom-oxygen bond distances of 1.54, 1.49 and 1.42 A respectively. The apparent rate constants for the oxygen exchange reactions of the potassium salts shown in Table 2 would seem to suggest that the oxygen exchange lability is related both to the charge of the central atom and to the central atom-oxygen bond distance, A similar trend is observed in the case of the oxygen exchange between water and these anions.t~8) There does not appear to be the same parallel correlation between the apparent rate constant and the thermodynamic or thermal stabilities of these salts as was observed with the other salts. For example, the standard heat of formation of potassium sulphate is much greater than that of potassium perchlorate t21) and the respective decomposition temperatures are > 1200° and 550°. The fact that the final decomposition products of potassium sulphate are potassium oxide and sulphur trioxide but those of potassium perchlorate are potassium chloride and oxygen is probably related to the differences in the relative stabilities of the salts. The mechanism o f the exchange reaction
The mechanism of the interaction of oxygen gas with the salt appears to be complex. It must involve at least three processes which probably occur at the surface of the salt:el) molecular or atomic adsorption of the gas, oxygen exchange with the anion and finally desorption of the gas. During the exchange process itself, interaction of the adsorbed oxygen with the salt may cause an increase in the co-ordination number of the salt which would allow for the oxygens to become equivalent. The adsorptiondesorption rates are not necessarily equal unless the adsorption is physical in nature. Studies of the adsorption of oxygen on the surfaces of numerous oxides show that the adsorption process is much faster than the desorption process ~1.2°-311 and that electron transfer takes place during the adsorption process which results in the production of various species such as O2-, Ozz-, O- and 0 3-. Thus, the adsorption and desorption processes probably have independent activation energies which are a function of many parameters such as surface geometry and lattice defects. Although the apparent activation energies reported here do show specific trends, they should not necessarily be taken as a measure of the relative strength of the central atom-oxygen, bond. ~6~ Ibid. p. 57. ~,7~R. S. LEEand G. B. CAaVErcrER,J. Phys. Chem. 63, 279 (1959). ~28jj. O. EDWARDS,Inorganic Reaction Mechanisms, p. 141. Benjamin,New York (1964). ~9~ T. SMrm, J. Electrochem. Soe. 111, 1020, 1027 (1964). ca0jW. E. GARNER,F. S. STONEand P. F. TruEr, Proc. Roy. Soc. A211, 472 (1952). ~81~ T . I. BARRYand F. S. STONE,Proc. Roy. Soc. A~5, 124 (1960).
Oxygen-18exchange reactions betweengaseous oxygenand certain salts
2169
A large variation in the apparent activation energy for the oxygen exchange process is noted. This variation is probably a measure of the differences in the nature of the surface and the adsorption and desorption processes for the various salts. Surface area measurements are usually carried out at liquid nitrogen temperature based on the assumption that one layer of gas is physically adsorbed. Such a surface is not necessarily related to the surface at elevated temperatures since many lattice defects could be produced or removed upon heating. In fact, the surface area has been shown by BARRY and STONEtal~ to vary markedly with the method of preparation and the heat treatment of the sample. A detailed interpretation of the exchange process itself requires a knowledge of the surface properties of each individual salt at the temperature of the exchange reaction since any variation in the nature of the surface probably influences the amount of oxygen adsorbed and subsequently the rates of adsorption and desorption. In conclusion, it is suggested that the apparent rate constant at any specific temperature appears to be a better measure of the central atom-oxygen bond lability than is a comparison of the apparent activation energies.
OXYGEN-18 EXCHANGE REACTIONS BETWEEN GASEOUS OXYGEN AND CERTAIN OXYGENCONTAINING INORGANIC SALTS* B. Z. SHAKHASHIRIand G. GORDON Department of Chemistry, University of Maryland, College Park, Maryland
(Received 27 January 1965) AImtraet--Rate measurements have been carried out on the 180 exchange reaction between oxygen gas and certain alkali metal chromates, molybdates, tungstates, phosphates and sulphates in the 600-850°C temperature range. It was observed that the rate of exchange was influenced by the nature of the cation and the nature of the central atom. In the alkali metal chromate series, the potassium salt exhibited the least tendency to exchange, but the activation energy for the exchange process was observed to decrease with an increase in the cation radius. In contrast, both the rate of exchange and the activation energy decreased as the effective nuclear charge of the central atom increased, as the central atom-oxygen bond distance decreased, and as the oxidation state of the central atom increased. A mechanism is proposed for the exchange reaction and a correlation between the rate of exchange and various thermal and thermodynamic properties is presented. INTRODUCTION
TH~ xsO exchange reactions between gaseous oxygen and solid oxides have been widely investigated by several workers. (x) These investigations have shown in general that exchange can take place if the temperature is above 300° and that the initially rapid exchange which does take place on the surface is followed by a slow exchange, whose rate is controlled by the diffusion of oxygen from the inner layers of the lattice to the surface. SPITSYN and FINIKOVt2) have studied the isotopic exchange between gaseous oxygen and the alkali metal sulphates in the 680-820 ° temperature range. They found that the exchange proceeds not only on the surface but also in the inner layers of the crystals. The activation energies for the exchange processes with lithium, sodium, potassium, rubidium and caesium sulphates were 39, 54, 57, 40 and 24 kcal/mole respectively and were based on the assumption that the limiting step occurred on the surface. It was shown that the activation energy depends on the nature of the cation and that marked variations in the relative strength of the oxygen bonds do exist. This study was undertaken in an attempt to determine the effect of various salt parameters, such as the effective nuclear charge of the central atom, the central atom-oxygen bond distance, and the oxidation state of the central atom on both the rate of exchange and the activation energy. * This work was supported by the United States Atomic Energy Commission under grant No. AT-(40-1)-2858 and is based on a thesis submitted by B. Z. S. to the Graduate School of the University of Maryland in partial fulfillment for the degree of Master of Science (1964). (t) For example, see E. R. S. WINTER, Advances in Catalysis, X, pp. 196-241. Academic, New York
(1958). (2) V. I. SPITSVNand V. G. FINIKOV, Dokl. Akad. Nauk SSSR. 108, 491 (1956). CA 51:3249h. 2161
2162
B . Z . SHAKHASHIRIand G. GORDON EXPERIMENTAL
General procedure The technique employed for studying the isotopic exchange between gaseous oxygen and the various salts consisted of placing a 30-50 mg sample of an ~80-enriched salt (0.60-1.23 ~ xsO) in a Vycor or quartz tube contained in a small furnace. The sample was degassed for a minimum of one hour (<10 -4 torr) at the temperature at which the exchange reaction was to be studied. Approximately 5000 torr-ml of oxygen gas of normal isotopic composition (0.20 ~ ~80) was introduced into the reaction vessel. A small capillary tubing connected the reaction vessel to the inlet system of a Nuclide Analysis Associates RMS-11 Isotope Ratio Mass Spectrometer and allowed a continuous flow of gas. The total volume occupied by the gas was about 100 ml. The progress of the reaction was monitored mass spectrometrically by periodically measuring the change in the ratio of the mass 34-32 (xnO180 to leO160) peaks. Samples of the enriched salts were heated for two hours at various temperatures from 500-850 ° in the exchange vessel which was open to the mass spectrometer. The extent of decomposition, if any, in the temperature range of the study of each salt was checked by comparing the mass spectra of the residual gases found in the vacuum line before and after the heat treatment and by observing any loss in weight of the sample.
Materials All the salts used in this study were enriched in 180. The method of enrichment depended on the nature of the oxy-anion, on the relative ease with which it exchanged oxygen with water, and on the solubility of the resulting compound. In all methods water from the Isomet Corporation which contained 1"55 % xsO was used. Chromium (VI) oxide. Ten grammes of anhydrous chromium (VI) oxide was dissolved in 18 ml of xsO-enriched water in a glass stoppered flask and allowed to stand for 1 hr, since it has been shown ~8~ that isotopic equilibrium is reached very quickly. Portions of the resulting oxygen-18 enriched chromate solution were used directly in the following preparations. Lithium chromate. Two and one half grammes of anhydrous lithium hydroxide was dissolved in 8 ml of distilled water. Excess ~80-enriched chromate solution was added to destroy any lithium carbonate that might have been present. In order to insure the presence of mononuclear chromium species, solid lithium hydroxide was added ¢4~ until the pH of the solution was 8. The solution was evaporated slowly in this and all of the following preparations in order to obtain the appropriate pure salt. Sodium chromate and potassium chromate. BALOGAand EARLEY15~report a half-life of 17 hr for the oxygen exchange reaction between the chromate ion and water at 25 °. Therefore, 17-5 g of anhydrous sodium chromate was dissolved in 20 ml of xsO-enriched water and allowed to stand for 6 days in a glass stoppered flask. Similarly, 11.6 g of anhydrous potassium chromate was dissolved in 20 mi of 1sO-enriched water and allowed to stand for 6 days in a glass stoppered flask. Rubidium chromate and caesium chromate. Twelve grammes of anhydrous rubidium carbonate and 20 g of anhydrous caesium carbonate was treated separately with stoichiometric amounts of the asO-enriched chromate solution. The final pH was 8. Sodium molybdate and sodium tungstate. HALL and ALEXANDEac~ observed that the exchange reactions of oxygen between water and both molybdate and tungstate ions were fast. Ten grammes of anhydrous sodium molybdate was dissolved in 18 ml of 180-enriched water and allowed to stand 10 days in a glass stoppered flask. The pH was adjusted to 8 by means of a carbonate-free sodium hydroxide solution in order to insure the presence of mononuclear species of molybdenum/v~ Similarly, 10 g of anhydrous sodium tungstate was dissolved in 18 ml of 1sO-enriched water and allowed to stand for a minimum of 10 days in a glass stoppered flask. The pH was adjusted to a ~s~ I. DOSTROVSKYand D. SAMUEL,Inorganic Isotopic Syntheses, R. H. HERBER (Editor), p. 139. Benjamin, New York (1962). la) y . I. SASKI, Acta Chem. Scand. 16, 719 (1962). ~ M. R. BALOGAand J. E. EARLEY,J. Phys. Chem. 67, 964 (1963). ~n~N. F. HALL and O. R. ALEXANDER,J. Amer. Chem. Soc. 62, 3455 (1940). tT) y . I. SASKI, I. LINDQVISTand L. G. SILLEN,J. Inorg. NucL Chem. 9, 93 (1959).
Oxygen-18 exchange reactions between gaseous oxygen and certain salts
2163
value of 10 by means of a carbonate-free sodium hydroxide solution in order to insure the presence of mononuclear species of tungsten.C8 Potassium phosphate. Fourteen grammes of anhydrous potassium dihydrogen phosphate was dissolved in 27 ml of 180-enriched water and the solution was heated under reflux at 90 ° for 6 days. The solution was cooled to room temperature and the pH was adjusted to 13 by means of a potassium hydroxide solution. Potassium sulphate. HALL and ALEXANDERt°) and WINTER and co-workers, ~9~ have shown that the sulphate ion and water do not exchange oxygen while the rate of oxygen exchange between sulphite ion and water was measurable. Therefore, 8 g of anhydrous potassium sulphite was dissolved in 14 ml of 180-enriched water and allowed to stand overnight in a glass stoppered flask. The solution was electrolysed by using a sheet of platinum of large surface area (about 6 cm ~) as the anode and a platinum gauze as the cathode for a period of at least 6 hr at a current density of about 0.3 A/cm ~. It was necessary to add a few milliliters of lsO-enricbed water to the cell to compensate for evaporation losses. The amount of sulphite left was determined to be less than 1 ~ by testing with a standard permanganate solution. The remaining sulphite was removed by boiling the solution after the addition of a few drops of concentrated sulphuric acid. Potassium perchlorate. HOERING and co-workers t1°~ have estimated the half-life for the oxygen exchange reaction between the perchlorate ion and water at room temperature to be greater than 100 years. They also observed that the rate of the oxygen exchange reaction between the chlorate ion and water is measurable at 100 °. Eleven grammes of anhydrous sodium chlorate was dissolved in 27 ml of 1sO-enriched water, to which 1.2 ml of concentrated perchloric acid was added. The solution was heated under reflux overnight in order to allow isotopic equilibrium to obtain and then it was electrolysed for a period of 7 hr. Analysis showed 99.9 ~ conversion of chlorate to perchlorate. The resulting solution was treated with concentrated potassium nitrate solution to precipitate the potassium perchlorate. In all cases, the freshly precipitated salt was collected by filtration, washed once with ethyl alcohol or acetone and once with ether. Each salt was dried in an oven for 2 hr at 400 ° and dried further by heating for an additional 2 hr under vacuum at a temperature well below the melting point or decomposition temperature. Oxygen and carbon dioxide of normal isotopic composition (0'20 ~ asO) were obtained from the Matheson Company.
Isotopic analysis of oxygen in the salts The oxygen isotopic analysis of the salts was carried out by converting the oxygen into a gaseous form which was more suitable for mass spectrometry. Since in general the decomposition of the salts takes place at temperatures in excess of 1000 °, the salts were reduced with mercury (II) cyanide to form carbon dioxide with the exception of potassium perchlorate and lithium chromate which were decomposed at 550 ° and 700 ° respectively. The method of reduction to carbon dioxide was that of ANBAR and GUTTMANN.t11) The samples were heated for 2 hr at 400 ° with a mixture of mercury (II) chloride and mercury (II) cyanide; the carbon dioxide formed was purified over zinc amalgam at 200 ° . The isotopic composition of the carbon dioxide was determined mass spectrometrically by measuring the mass 46/(44 + 45) ratio (12ClaOXsO/lzCaeOa60 + 13C160160). The contribution of the x~O isotope is negligible for these samples which contain less than 1-55 ~ 180. The enrichment of the salt was calculated from the measured ratio for the enriched and normal samples.
Apparatus The exchange vessel consisted of a Vycor or a quartz tube 22 m m × 21 cm equipped with a 24/40 standard taper joint. The salt sample was weighed in a small Vycor tube (13 m m × 2 cm) and was placed in the exchange vessel. The vessel was evacuated on a high vacuum line which was equipped with a two stage mercury diffusion pump and a rotary type oil pump. The pressure was monitored ts~ y . I. SASKI, Acta Chem. Scand. 15, 175 (1961). tg~ E. R. S. WINTER, M. CARLTON and H. V. A. BRISCOE, J. Chem. Soc. 131 (1937). c10~T. C. HOERING, F. T. ISH/MORI and H. O. McDONALD, J. Amer. Chem. Soc. 80, 3876 (1958). iltl M. ANBAR and S. GUTrMANN, Int. J. Appl. Rad. Isotopes 5, 233 (1959).
2164
B.Z. SHAKHASHIRIand G. GORDON
by means of RCA 1946 Thermocouple Gauge. The sample was heated by a small furnace which was constructed in the following manner: Chromel "A" resistance wire was wound about the vessel, covered with asbestos and gypsum, and connected to a Variac; the resistance of the furnace varied from 15 to 29 f~ and the total length of the furnace was about 14 cm. A more elaborate furnace was constructed as follows: chromel "A" resistance wire (80 f~, 150 W at 110 V) was wound about the vessel and covered as before and connected to a Variac through a temperature regulator, two wires of lower resistance were connected in parallel (46 t), 270 W at 110 V) to another Variac. It was found that the second type furnace reached the desired temperature much more quickly than the first type, but the variation in equilibrium temperature of both furnaces was less than 4-0.5 °. The actual temperature control was maintained by means of a Model 283C Barber-Coleman Temperature Regulator. The temperature was measured by means of an insulated chromel-alumel thermocouple obtained from Advanced Technology Laboratories which was placed through a small Vycor tubing which extended into the middle of the exchange vessel. The output of the thermocouple was determined with a Calibration Standards Corporation VA-100A Precision Volt-Amp Meter to a precision of +0.01 mV. RESULTS Vycor and quartz were shown not to exchange oxygen with gaseous oxygen at 950 °. F o r example, after a period o f two hours, the change in the enrichment o f an 8 fold enriched (12) sample o f gaseous oxygen was less than 0.3 ~o. This is in agreement with HUTCHINSON (13) who reports no detectable exchange in the silica-oxygen system below 1000 °. The results o f individual exchange experiments were analysed by means o f the McKAY equation, c14a5) In (1 -- Ft) = --[(a + b)/(ab)](Rt)
(1)
where a and b are the n u m b e r o f moles o f exchangeable oxygen contained in the salt and gas respectively, R is the rate o f exchange in mole sec -1, t is the time in seconds and F t is defined as the fraction F , = ( N t - - No)/(N~o - - N t )
(2)
at any time t. The quantity N t is the measured 180 to 160 ratio o f the gas at any time t, N Ois the initial ratio, and No~ is the ratio at isotopic equilibrium. No~ was calculated according to Equation (3) Noo =
n,(No)i
ni
(3)
0
where i is the n u m b e r of oxygen-containing components in the system and n is the n u m b e r o f moles o f oxygen. The validity o f Equation (3) was checked experimentally. The agreement between the calculated and measured ratio was better than 0.4 ~o (for example, calc. 4.98 fold; measured, 5.00 fold). The calculation o f the rate o f exchange requires an exact knowledge o f the surface area o f the salt and the n u m b e r o f moles o f oxygen in the gas phase. The ideal gas law was used to calculate the moles o f gas in the gas phase; however, the surface areas o f the salts at the temperatures o f the reactions are not known. Therefore, an apparent (12)Oxygen of normal isotopic composition would correspond to an enrichment of 1-00 on this scale. t18) D. A. HUTCHINSON,J. Chem. Phys. 22, 758 (1954). (14) H. A. C. McKAY, Nature, Lond. 142, 997 (1938). (~5)All of the calculations were done on the I.B.M. 7094 Computer in the University of Maryland Computer Science Center.
Oxygen-18 exchange reactions between gaseous oxygen and certain salts
2165
rate constant, k, was defined as [R(a + blab)] and was obtained directly from the slope of the McKAy plot. t14) Similarly, an apparent activation energy, E~, was obtained from the slope of a graph of log k as a function of the reciprocal of the absolute temperature. Although neither k nor E , are absolute quantities, they can be useful for the purpose of comparison of the effect of variation in the salt parameters discussed earlier.
I'0 0'9 0'8 0'7 0"6
0'5 I-F 0'4
0,3
0"2
I 600
I 1200
I I 1800 2400 T I M E (SECONDS)
I 3000
I 3600
I 4,¢00
FIG. 1.--Graph of log (1 -- F) as a function of time for the K~CrO~(s)~O2(g) exchange reaction at 855°. The results of a typical experiment are shown in Fig. 1. Presumably, the fast rate of exchange is that taking place at the surface of the salt while the slower rate is controlled by diffusion from the inner layers of the salt to the surface. This is in agreement with the studies of WINTERtl~ on the reaction between gaseous oxygen and metal oxides and the studies of SPITSYN and FINIKOVt~) who report that the exchange process between gaseous oxygen and the alkali metal sulphates proceeds not only on the surface, but also takes place in the inner layers of the salt. In this study the extent of the exchange reaction was followed for values of F < 0.35 which corresponds to the linear region of the McKAY plot. The individual values of k were reproducible to better than 3 per cent and the standard deviation in the apparent activation energy was no greater than q-3 kcal/mole, except in the case of sodium chromate where the deviation was + 10 kcal/mole. The apparent rate constants and apparent activation energies for the oxygen exchange reaction between gaseous oxygen and the alkali metal chromates are shown in Table 1. These data were used to calculate apparent rate constants at other temperatures and those at 623 ° are also recorded in Table 1.
2166
I.
Z.
TABLE 1 . - - T H E EFFECT
SHArd~ASHI~ and G. GORDON OF THE CATIONON THEHETEROGENEOUSOXYGEN EXCHANGE REACTION
Salt
(sec-1)
e2a ° X 105 kenle" (sec-t)
Li2CrO4 Na2CrO4 K2CrO~ Rb2CrO4 Cs2CrO~
6"64* 6"86t 1"58~ 12"5§ 7"89¶
6"64 0.007 0.004 1"88 1.69
kmeas • >( 1 0 5
• 623°
~ 712°
~ 741 °
Temp. Range E~ (°C) (kcal/mole) 598-623 680-712 723-741 713-745 675-710
§ 713°
152 135 92 37 31
¶ 710°
The alkali metal chromates were observed to change colour f r o m yellow to red when heated above 150 °. The colour change was reversible except in the case o f lithium chromate which developed a deep b r o w n colour u p o n cooling to r o o m temperature. The results o f the oxygen exchange reaction between gaseous oxygen and various other sodium and potassium salts are shown in Table 2. TABLE 2.--THE EFFECT OF THE CENTRAL ATOM ON THE HETEROGENEOUS OXYGEN EXCHANGE REACTION 778 ~ keale ' x 105 (sec-a)
kmeas" X 10 5
Salt Na2CrO4 NazMoO4 Na2WO4 K3PO4 K~SO4 KC10~ * 712°
(sec-1)
6"86* 521 l'50t 1-50 0.57:~ 0"18 5"10 § 21"5 0"52¶ 0"97 decomposed without exchange[I ~ 778°
~ 808°
§ 740°
Temp. Range E~ (°C) (kcal/mole) 680-712 746-810 808-845 727-767 741-851 ¶ 741°
135 94 85 80 36
Jl 550°
DISCUSSION The effect o f the cation The alkali metal chromates and sulphates (except for lithium and sodium) have been observed to be isomorphous (16-1s) and each c h r o m i u m a t o m is tetrahedrally surrounded by four oxygen atoms at an average distance of 1.60 A. I f the apparent rate constants can be used as a measure o f the order o f lability with respect to oxygen exchange, then the order for the alkali metal chromates at 623 ° is lithium > rubidium ~ caesium > sodium > potassium as shown in Table 1. This is in agreement with the observations o f SPITSYN and FINIKOV. t2} At 710 ° the same relative order obtains (neglecting lithium chromate which would decompose at 700°). Even t h o u g h the apparent activation energy m a y change with large variations in temperature, it is interesting to note the apparent rate constants at 740 ° where the lability o f sodium chromate becomes nearly equal to that o f rubidium chromate. j. W. MELLOR, A Comprehensive Treatise on Inorganic and Theoretical Chemistry, Vol. XI, pp. 245, 251. Longmans, Green, New York (1931). (xT~H. W. SMITHand M. Y. COLBY,Z. Krist. 103, 90 (1941). (18~j. j. MILLER,Z. Krist. 99, 32 (1938). lle)
Oxygen-18exchangereactionsbetweengaseousoxygenand certain salts
2167
It is generally accepted that the polarizing power of ions increases with increasing positive charge but decreases with increasing size. Observations that salts of oxyanions with large cations tend to be more stable thermodynamically than salts of the same oxy-anion and small cations suggest that the difference may be associated with the polarization of the anion by the cation, ttg) Therefore, as the degree of polarization increases, the stability of the anion should be expected to decrease. On the other hand, SPITSYN(~) has suggested the possibility of mutual polarization between all of the atoms which form the oxy-anion salt. The relatively strongly polarizing lithium and sodium ions would be slightly polarized by the oxygens of the anions, but the relatively weakly polarizing rubidium and caesium ions would be more polarized by neighbouring oxygens, which would result in a weakening of the central atom-oxygen bond. Since these effects oppose each other, it is indeed possible that potassium salts could show the least tendency to undergo oxygen exchange if the extent of these interactions were maximized. This interpretation would be in agreement with other results tz~ which suggest that the potassium salt does exhibit the least tendency to exchange. The standard heats of formation of the alkali metal sulphates increase with potassium > lithium > rubidium > caesium > sodium, t~l~ This order is markedly different from that observed for the oxygen exchange labilities. The effect of the bond distance The central atom-oxygen bond distance in the isomorphous potassium chromate and sulphate t ~ are 1-60 and 1.49 A respectively. The apparent rate constants shown in Tables 1 and 2 indicate that the chromate anion is more labile with respect to oxygen exchange than is the sulphate anion. This is in agreement with the thermodynamic stability of these salts in that the standard heat of formation of potassium sulphate is larger than that of potassium chromate.t2~ The effective nuclear charges of the central atoms of these anions differ only slightly even though the atomic numbers of sulphur and chromium are 16 and 24 respectively. However, sulphate and chromate differ markedly in their central atom-oxygen bond distances. This suggests that perhaps the apparent rate constant for the heterogeneous oxygen exchange reaction can be used as a measure of the relative strength of the central atom-oxygen bond. It is interesting to note that the rate of oxygen exchange between water and the two anions follows the same trend. ~5'6~ The effect of effective nuclear charge of the central atom The apparent rate constants shown in Table 2 suggest that the order of lability with respect to oxygen exchange is chromate > molybdate > tungstate which is the same as the order of decreasing thermodynamic stability.C21~ The observations of SeITSYN and KULESHOV{~°'~ indicate that sodium tungstate is also more stable thermally than the isomorphous sodium molybdatet24'~5~ in that the loss in weight t19~R. T. SAI'a~RSON,ChemicalPeriodicity, pp. 162-166. Reinhold,New York (1960). t~o~V. I. SPITSX,N, Zhur. Obsch. Khim. 17, 11 (1947); 20, 550 (1950); CA 44:6243c. t21}W. M. LATIMER,Oxidation Potentials (2nd ed.), pp. 329-335. Prentice-Hall,New York (1952). {~ M. A'ro~tand R. E. Rtn,rDLE,J. Chem. Phys. 29, 1306 (1958). t~a~V. I. SerrsvN and I. M. KUL~HOV,Zhur. Obsch. Khim. 21, 401 (1951); CA 45:5553i. ~24}I. LINDQVIST, hcta Chem. Scand. 4, 1066 (1950). 12s~A. F. WELLS, Structural Inorganic Chemistry (3rd ed.), p. 445. Clarendon, Oxford (1962). 3
2168
B.Z. SrIAKHASrmUand G. GORDON
after heating for 6 hr at 1200° was 2.30 ~o for the molybdate and 0.59 ~o for the tungstate. Molybdenum and tungsten, are tetrahedrally surrounded by four oxygen atoms at an average distance of 2.0 A, but the chromium-oxygen bond distance is 1"60 A. Apparently, the increase in the effective nuclear charge of the central atoms has a greater influence than the increase in central atom-oxygen bond distance in that sodium tungstate is more stable thermodynamically, thermally and with respect to oxygen exchange than the corresponding molybdate and chromate. The effect o f the oxidation state o f the central atom
The isoelectronic phosphate, sulphate and perchlorate anions are tetrahedralt22'26,zT~ with central atom-oxygen bond distances of 1.54, 1.49 and 1.42 A respectively. The apparent rate constants for the oxygen exchange reactions of the potassium salts shown in Table 2 would seem to suggest that the oxygen exchange lability is related both to the charge of the central atom and to the central atom-oxygen bond distance, A similar trend is observed in the case of the oxygen exchange between water and these anions.t~8) There does not appear to be the same parallel correlation between the apparent rate constant and the thermodynamic or thermal stabilities of these salts as was observed with the other salts. For example, the standard heat of formation of potassium sulphate is much greater than that of potassium perchlorate t21) and the respective decomposition temperatures are > 1200° and 550°. The fact that the final decomposition products of potassium sulphate are potassium oxide and sulphur trioxide but those of potassium perchlorate are potassium chloride and oxygen is probably related to the differences in the relative stabilities of the salts. The mechanism o f the exchange reaction
The mechanism of the interaction of oxygen gas with the salt appears to be complex. It must involve at least three processes which probably occur at the surface of the salt:el) molecular or atomic adsorption of the gas, oxygen exchange with the anion and finally desorption of the gas. During the exchange process itself, interaction of the adsorbed oxygen with the salt may cause an increase in the co-ordination number of the salt which would allow for the oxygens to become equivalent. The adsorptiondesorption rates are not necessarily equal unless the adsorption is physical in nature. Studies of the adsorption of oxygen on the surfaces of numerous oxides show that the adsorption process is much faster than the desorption process ~1.2°-311 and that electron transfer takes place during the adsorption process which results in the production of various species such as O2-, Ozz-, O- and 0 3-. Thus, the adsorption and desorption processes probably have independent activation energies which are a function of many parameters such as surface geometry and lattice defects. Although the apparent activation energies reported here do show specific trends, they should not necessarily be taken as a measure of the relative strength of the central atom-oxygen, bond. ~6~ Ibid. p. 57. ~,7~R. S. LEEand G. B. CAaVErcrER,J. Phys. Chem. 63, 279 (1959). ~28jj. O. EDWARDS,Inorganic Reaction Mechanisms, p. 141. Benjamin,New York (1964). ~9~ T. SMrm, J. Electrochem. Soe. 111, 1020, 1027 (1964). ca0jW. E. GARNER,F. S. STONEand P. F. TruEr, Proc. Roy. Soc. A211, 472 (1952). ~81~ T . I. BARRYand F. S. STONE,Proc. Roy. Soc. A~5, 124 (1960).
Oxygen-18exchange reactions betweengaseous oxygenand certain salts
2169
A large variation in the apparent activation energy for the oxygen exchange process is noted. This variation is probably a measure of the differences in the nature of the surface and the adsorption and desorption processes for the various salts. Surface area measurements are usually carried out at liquid nitrogen temperature based on the assumption that one layer of gas is physically adsorbed. Such a surface is not necessarily related to the surface at elevated temperatures since many lattice defects could be produced or removed upon heating. In fact, the surface area has been shown by BARRY and STONEtal~ to vary markedly with the method of preparation and the heat treatment of the sample. A detailed interpretation of the exchange process itself requires a knowledge of the surface properties of each individual salt at the temperature of the exchange reaction since any variation in the nature of the surface probably influences the amount of oxygen adsorbed and subsequently the rates of adsorption and desorption. In conclusion, it is suggested that the apparent rate constant at any specific temperature appears to be a better measure of the central atom-oxygen bond lability than is a comparison of the apparent activation energies.