ONTARIO CHEMISTS. Dr. Howard Clark, Dr. Keith Dixon, and W. J . Jacobs (left to right) of the University of Western Ontario, London, Ont., study spectra which characterize products f r o m the reaction of trans-chlorohydribis(triethylphosphine)p l a t i n u m ( l l ) and tetrafluoroethylene in a sealed borosilicate glass tube
acid, and carbon monoxide in benzene in a stainless steel autoclave. X-ray work shows that the pentafluorosilicate and the fluoroborate are not isomorphous, although their unit cells are related. Infrared, nuclear magnetic resonance, and mass spectroscopic studies were also used to characterize the reaction products. A number of control reactions were carried out in a stainless steel autoclave to pinpoint the path. T h e carbonyl complex is of considerable interest, Dr. Clark notes, since it is isoelectronic with transchlorocarbonylbis(triphenylphosphine)i r i d i u m ( I ) , known as Vaska's complex. Dr. James P. Collman of the University of North Carolina has shown Vaska's complex yields a stable iridiumnitrogen complex in reactions with organic azides (C&EN, Jan. 16, page 42). Dr. Clark observes that the platinum carbonyl ion is m u c h less reactive than Vaska's complex or its rhodium analog. It does not combine with hydrogen or dry hydrogen chloride at atmospheric pressure at 2 5 ° C. This is consistent with the increase in promotional energy required to form two more bonds caused by t h e decrease in electron density on the metal atom. Also, unlike Vaska's complex, there is no reaction with benzoyl azide in chloroform solution. Dr. Clark does find that the carbonyl is readily displaced by chloride ion. It will b e interesting to see, he says, how variations in the phosphine and halide ligands affect the chemistry of this remarkably generated platinum carbonyl.
Monomeric SiCI2 reacts with chlorides 154TH
ACS NATIONAL MEETING Inorganic Chemistry
Monomeric silicon dichloride reacts with carbon tetrachloride, boron trichloride, and phosphorus trichloride to form, respectively, SiCl 3 CCl 3 , SiCl 3 BC1 2 , and SiCl 3 PCl 2 . But no simple compound seems to result from silicon dichloride and stannic chloride, according to Dr. Peter Timms of the University of California, Berkeley. From structural studies of the reaction products, Dr. Timms concludes that SiCl 2 inserts readily into M-Cl bonds ( M = B, P, C ) . This offers a new synthetic route from M-Cl to M-SiCl 3 , he says. T h e new compound SiCl 3 PCl 2 is a little less volatile than Si 2 Cl 6 , and melts at —64° C. It is unstable at 25° C , decomposing to SiCl 4 and yellow PCI polymers. Silicon dichloride is a carbenelike molecule (singlet ground s t a t e ) , potentially reactive toward many compounds. Chemists have known for years that it can be prepared from silicon and silicon tetrachloride at about 1350° C. It is extremely shortlived, polymerizing to ( S i C l 2 ) n or reacting with excess silicon tetrachloride to form higher perchlorosilanes. If p r e p a r e d at low enough pressure, silicon dichloride can b e transferred to a liquid nitrogen-cooled surface without appreciable gas-phase polymeriza-
tion or reaction with silicon tetrachloride. In Dr. Timms' technique, vapors of various compounds are injected into a stream of SiCl 2 monomer emanating from a furnace and t h e products cocondensed on an adjacent surface at —196° C. T h e apparatus is identical to that which Dr. Timms used for cocondensations with boron monofluoride (C&EN, S e p t 19, 1966, page 5 0 ) . At 1350° C , Dr. Timms and his coworkers prepare silicon dichloride in better than 9 5 % yield. They keep the pressure of permanent gas below 5 X 10" 6 mm. H g by fast pumping. T h e resulting condensate is warmed to room temperature to separate volatile products from any polymer that may have formed. The volatiles are then separated on a low-pressure, low-temperature distillation column. Isolable volatiles are characterized by mass spectrum and infrared absorption analyses. T h e reaction of SiCl 2 with BC1 3 provides a good example of Dr. Timms' procedure. At first, a blue solid forms on the cold surface. W h e n w a r m e d to room temperature, this solid evolves a mixture of silicon tetrachloride, higher perchlorosilanes, and a liquid slightly more volatile than Si 2 Cl 6 . T h e mass spectrum of this liquid and its gas-phase infrared spectrum support a formula SiCl 3 BCl 2 with a silicon-boron bond. T h e compound melts at — 72° C. and is stable at room temperature in the absence of air. Dr. Timms previously m a d e the same compound by passing boron trichloride over silicon, or silicon tetrachloride over boron, at 1400° to 1700° C , using the same apparatus. Dr. A. G. Massey and Dr. D . S. A. Urch of Queen Mary College, London, also prepared it in tiny amounts, in 1964, using a mercury discharge in BC1 3 in a quartz cell. T h e reaction of BC1 8 with SiCl 2 is unlike that of B F 3 with SiF 2 . (At 1 micron Hg, SiF 2 is stable for several minutes and does not react with SiF 4 . By contrast, the lifetime of SiCl 2 may b e only a few milliseconds.) The B F 3 - S I F 2 reaction yields SiF 3 SiF 2 BF 2 and its higher homologs, but no SiF 3 BF 2 . T h e compound SiCl 8 SiCl 2 BCl 2 did not form in the BCl 3 -SiCl 2 reaction. Thus, SiCl 2 inserts into the B-Cl bond, Dr. Timms postulates, and SiF 2 initially forms a diradical *Si2F4% which then reacts with other molecules. However, SiCl 2 reacts similarly to SiF 2 with benzene and acetylene. These reactions suggest that in nonchlorinated systems, where insertion into a bond is not favored, SiCl 2 may, like SiF 2 , polymerize via reactive diradicals. SEPT. 18, 1967 C&EN
57
Dr. Timms has also produced SiF 3 B F 2 in low yields (about 5%) by re acting SiCl 2 with B F 3 . In this reac tion, SiCl 2 FBF 2 probably forms first. It then exchanges halogens (with other SiCl 2 FBF 2 molecules ) to give the fully fluorinated compound. Recently, Dr. Timms has succeeded in preparing monomeric silicon dibromide in 9 0 % yield from silicon and silicon tetrabromide. T h e dibromide seems less reactive than the dichloride, Dr. Timms says, b u t it does react with boron trifluoride rather inefficiently to give S i F 3 B F 2 through the intermediate SiBr 2 FBF 2 . Although SiCl 2 has been known for 50 years, he notes, no one has pre viously a t t e m p t e d to apply high-tem perature, high-vacuum methods to its formation. And these are t h e only conditions under which SiCl 2 is a chemically interesting species, he adds.
Loyola chemist prepares and isolates Xe0 2 F 2 Elusive xenon dioxide difluoride, X e 0 2 F 2 , has been prepared and iso lated by Dr. John L. Huston of Loyola University ( C h i c a g o ) . During a mass spectrographic study of the chemical properties of xenon compounds, the Loyola chemist has prepared the com p o u n d by reacting X e O s and X e O F 4 [/. Phys. Chem., 7 1 , 3339 ( 1 9 6 7 ) ] . Earlier, scientists at Argonne Na tional Laboratory observed mass spec tra of X e 0 2 F 2 and X e O F 4 . T h e two compounds were formed from the step wise hydrolysis of X e F 6 , according to the Argonne workers. Since then, Argonne's Dr. C. L. Chernick, Dr. H . H . Claassen, J. G. Malm, and P. L. Plurien isolated X e O F 4 and studied its properties. But attempts by Dr. M. H . Studier of Argonne and Dr. Huston to prepare
Xe0 2 F 2 has distorted tetrahedral structure
and isolate X e 0 2 F 2 in substantial quantity by partial hydrolysis of X e F 6 or by partial reaction of X e F 6 with glass have not been successful. After preparing X e 0 2 F 2 by the re action of X e O s with X e O F 4 , Dr. Hus ton used a time-of-flight mass spec trometer to follow the purification of the X e 0 2 F 2 . X e O F 4 was easily re moved because of its volatility. H o w ever, X e F 2 was more difficult to remove from X e 0 2 F 2 . As X e F 2 was removed, those portions of the mass spectrum due to X e 0 2 F 2 alone were enhanced compared to those portions due to both X e 0 2 F 2 and XeF2. Both positive and negative ion mass spectra of X e 0 2 F 2 were taken. T h e negative ion mass spectrum consisted of XeF-, X e F 2 - , and XeOF~. T h e dioxide difluoride can decom pose to X e F 2 and 0 2 , Dr. Huston finds. T h e decomposition rate depends on earlier conditioning of the containing apparatus. X e Q 2 F 2 reacts with X e F 6 in a rather slow reaction to produce X e O F 4 . Dr. Huston has also observed that X e O F 4 forms in Kel-F apparatus overconditioned by XeF 6 . T h e Loyola chemist has found that when X e 0 2 F 2 is exposed to moist air, it rapidly hydrolyzes to X e 0 3 . But he has also detected a faint ozonelike odor resembling that of xenon tetrox ide. Dr. Huston has confirmed the identity of X e 0 2 F 2 by distilling it into a silica container and analyzing the products formed by heating it to 300° C. H e and Dr. Claassen have confirmed the compound's identity by preliminary studies of the Raman and infrared spectra of the solid. The spectra are consistent with the pseudotrigonal bipyramidal structure for X e 0 2 F 2 predicted b y Dr. J. R. Gil lespie of McMaster University, Hamil ton, Ont. At Argonne, Dr. Claassen and sev eral others are still studying the com pound's structure, using a variety of techniques in addition to Raman and IR spectrometry. These techniques in clude x-ray diffraction, nuclear mag netic resonance, and Mossbauer spec trometry.
Myeloperoxidase breaks into 10 proteins
154TH
ACS NATIONAL MEETING Biological Chemistry
Myeloperoxidase, one of biology's most peculiar enzymes because of the strange behavior of one of its h e m e prosthetic groups, is showing further eccentricities. A group headed by Dr. Julius Schultz at H a h n e m a n n Medi cal College in Philadelphia finds 58
C&EN SEPT. 18, 1967
that the enzyme breaks down into 10 protein components under free flow electrophoresis in 6M urea. Further more, these 10 components divide into two groups on the basis of light absorption in the Soret region, where hemes are classically studied. Dr. Schultz and graduate students Norman Felberg and Sosamma John found that the first group, which migrated to the anode, was composed of four proteins with absorption maxima at 407 to 4 1 1 mμ. E a c h of the quartet of proteins showed weak enzymic activity and displayed a bleaching reaction on reduction with hydrosulfite. Hemes don't usually react this way on reduction; their Soret maxima usually tend toward the longer wave lengths. In this case, the green color of these four components dis appeared. T h e bleaching effect has caused the H a h n e m a n n scientists to wonder whether the prosthetic group on the first four components is in fact a conventional heme. T h e question, then, is exactly what type of maverick it is. T h e other six components, which were either neutral or migrated to the cathode, h a d strong enzymic activity and shifted from the typical between 424 and 428 mμ to 475 mμ on reduc tion, as myeloperoxidase classically does, Dr. Schultz says. Another feature of myeloperoxidase is its stability to 6M urea. Other enzymes lose their activity on such treatment, apparently through loss of tertiary structure. Yet, all six com ponents retained activity after electro phoresis in the urea medium. Dr. Schultz's group, therefore, isn't dis couraging speculation that the active six are isoenzymes and that the first four are denatured. T h e physiological role of myelo peroxidase is not at all understood. It is normally present in the lysosomal fractions of the types of white blood cells known as granulocytes. Myelo peroxidase is probably released during phagocytosis (when the white blood cell envelops and digests an invading bacterial cell). It may then possibly react with hydrogen peroxide at the bacterial membrane for eventual lysis. This action hasn't been shown in vivo, b u t test tube work by Dr. Warren Evans of the National Cancer Institute ( among others ) indicates that such action does take place outside the cell. T h e enzyme, according to Dr. Evans, is also useful as a routine diagnostic tool for the detection of that type of leukemia that affects the granulocytes. Detection of higher than normal concentrations of myelo peroxidase is enough to arouse sus picion that granulocytic leukemia is present.
