in this case, the ratio equals 1.

• Lithium is bound more strongly in the transition state than in the re-actant; ratio is less than 1.

• Lithium is bonded covalently to carbon in the original reactant, but not bound at all in the transition state; ratio equals 1.071.

The size of the observed kinetic iso­tope effect, 1.029, shows that the first two mechanisms aren't very likely, Dr. West says. But the value is close to the third possibility. Conclusion: Organolithium compounds are prob­ably covalent rather than ionic in their reactions in hydrocarbon solution. However, he points out, 1.029 is about halfway between 1 and 1.071, a value for a transition state with no lithium bonding; this means the reaction's ac­tual transition state probably includes some bonding to lithium.

Studies of infrared spectra also point to a covalent bond character for organolithium compounds, the re­search team says. Dr. West and Dr. Glaze ran vapor spectra on ethyl-lithhim, came across isotopic shifts in the low frequency region (600 to 300 cm,"1) upon substituting LiG. These shifts show that bands falling in this region come from carbon-lithium modes of vibration, Dr. West says. Lack of shifts in the higher frequencies (about 1000 cm. ]) means these bands aren't related to vibrations involving lithium to any great extent, he adds.

Unexpectedly low frequencies of the carbon-lithium modes may also indi­cate that the organolithium com­pounds associate highly into various polymeric structures in all states. Thus, these bands may not be simple C-Li stretch and deformation modes, could correspond to complex modes of the polymeric structures in which mo­tion of the lithium atoms plays a big role, Dr. West says.

Stable Complexes. The solid com­plexes were discovered by accident, Dr. West says. In one phase of the research on organolithiums, Dr. Glaze was making n-butyllithium by react­ing lithium and n-butyl bromide. After the reaction had apparently gone to completion, he filtered the mixture to remove unreacted halide, then let the clear solution stand. A day later, Dr. Glaze found that a white, crystal­line precipitate had formed.

Preliminary tests showed that the solid was not simply more unreacted lithium bromide, and that it contained equal amounts of base and halide.

Further study, Dr. West says, shows that the solid must be a complex with this composition: RLi nLiX (n equals 1 to 6 or more).

How did it form? Apparently, Dr. West reasons, not all of the alkyl halide reacts with the lithium and this sequence of reactions takes place:

• Li and RX give RLi plus LiX (solid).

• RX with excess RLi give RR and LiX, also a solid.

•L iX and RLi yield LiRnLiX i solid).

Dr. West and Dr. Glaze now use the reaction between an alkyllithium com­pound and an alkyl halide in hydro­carbon solution to make the complex. Previously, it was thought that the reaction gives lithium halide as a prod­uct, Dr. West says. But the lithium halide that precipitates carries with it excess alkyllithium compound in the form of the complex, he points out.

So far, complexes with a number of alkyl groups (ethyl, n-butyl, and cyclohexyl) have been used, Dr. West says. Halides are iodide and bromide. But, he notes, they can't make chloride or fluoride complexes.) Yields range from 75r/r to 90% based on the alkyl halide. This method of making the complexes, though, is in­direct and takes hours, he says.

Stable in Air. In contrast with either solid or liquid alkyllithium com­pounds, the complexes are stable in air. And they don't react readily with other chemical reagents. For instance, alkyllithium compounds react very quickly with ketones in a hydrocarbon solvent. But by complexing them, they can be exposed to Michler's ke­tone (/;,//-bis-dimethylaminobenzo-phenone) for 30 hours in a hydrocar­bon solution without any apparent re­action, Dr. West says. But adding ether destroys the complex, and the organolithium compound then reacts with the ketone immediately, he adds. Alkyllithium compounds can thus be deactivated by converting them to their complexes, then reactivated with ether.

The research workers are now look­ing for a better route to these com­pounds. With a more convenient, lower cost synthesis, Dr. West and Dr. Glaze see three possible uses for the alkyllithium-lithium halide com­plexes: as research chemicals, as in­dustrial reagents for making fine chemicals, and as catalysts for poly­merizing olefins.

Tin Oxidizes in Three Separate Steps The path of tin oxidation has been tracked with the help of electron microscope and electron diffraction techniques. Research at U.S. Steel's Applied Research Laboratory shows that the oxidation involves three sep­arate steps:

• Nucleation and growth of oxide platelets mark the start of oxidation.

• The platelets develop cavities which retard oxidation.

• Stress then causes cracks in the oxide film; more tin is exposed directly to oxygen and the oxidation rate speeds up.

Behind the research: efforts to im­prove corrosion (oxidation) resistance of tin plate under varied conditions. Such resistance will give longer storage life to tin plate used in can making, may also lead to improved, tin plated baking pans.

Using a plastic film stripping tech­nique to isolate tin oxide particles for electron microscope studies is the key, G. E. Pellissier told the 118th meeting of the Electrochemical Society in Houston, Tex. The technique also permits detailed study of oxide film topography.

Mr. Pellissier and co-workers W. E. Boggs and P. S. Trozzo first coat the oxidized tin surface with a thin film of a polyvinyl formal resin (Formvar) in p-dioxane. A thicker reinforcing layer of shredded, pure, and concentrated collodion (Parlo-dion) goes on top. Then a support­ing metal screen is pressed into the collodion layer. Mercury amalgama­tion removes the tin base, and a selec­tive solvent (like amyl acetate) re­moves the plastic. Electron micro­scope studies are then made.

Growth Centers. Mr. Pellissier and co-workers believe that tin oxidation usually begins with growth centers at dislocations in the metal crystal sur­face. Initially, these growth centers are too small to be seen microscop­ically. When the centers reach a diameter of about 1 micron, platelets of tin oxide begin to grow from them. The number of growth centers is about 107 per square centimeter, or roughly the same as the number of etch pits found on electropolished and vacuum annealed tin foil. This similarity tends to confirm that dislocations are the sites for growth nucleation.

During growth, a central "hub" ap-

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pears first; "spokes" radiate from the hub in a manner suggesting spiral growth. When the spokes are about 0.25 micron long, the space between them begins to fill, forming a rim about 1 micron in diameter. Soon, spikes grow laterally from the rim. Then, lath-shaped layers form from the spikes to make the larger single crystal plate­lets of oxide, Mr. Pellissier says.

The platelets cover much of the tin's surface by the end of the first stage and lower the oxidation rate. The platelets get thicker while the rate slows according to a logarithmic rate law. At the same time, cavities de­velop in the platelets at the oxide-metal interface. These cavities reduce the oxide area through which oxygen and tin ions diffuse, Mr. Boggs points out, further reducing the rate.

As the oxide film thickens, the cavities continue to grow. Stresses build up in the thick film and cause cracking, allowing direct access of oxygen to the unoxidized tin. At this point, oxidation increases generally but at an erratic rate, the scientists find.

Mr. Pellissier and co-workers iden­tify the oxide as the a-form of stan­nous oxide (tetragonal crystal struc­ture) from electron diffraction pat­terns. The crystal structure is always the same over a temperature range of 75° to 220° C. and 10 4 mm. to 500 mm. Hg oxygen pressure (conditions used in their experiment). At the lower temperatures, lengthy oxygen exposure is needed to get enough oxide for study. But at the lowest tempera­ture, no change from crystalline to amorphous forms occurs, contrary to previous findings. Mr. Pellissier says.

At oxygen pressures below 1.0 mm. Hg, the oxide film's microstructure changes radically but the crystal struc­ture remains the same. At such low pressures, the oxide forms mainly as dendritic crystals or filaments oriented along some crystallographic direction in each tin grain, Mr. Pellissier says. As oxygen pressure is increased to 1 mm. Hg, the filamentary crystals fill in to form more compact platelets.

Moreover, at oxygen pressures be­low 1 mm. Hg, the rate increases continuously over experimental pe­riods up to 3000 minutes, Mr. Boggs says. Under these conditions, the ox­ide film is discontinuous or spongy and provides no barrier between gaseous oxygen and tin, he concludes. Oxygen dissociation at the metal's surface is the oxidation rate's controlling factor in the low pressure region.