J. G. Ameen and H. F. Durfee
IBM-Components Division Endicott, New York 13760
I
I
The Structure of Metal Carbonyls
Since their discovery, metal carbonyls have been known in the chemical world novelty as zero valence state compounds. Up until ten years ago, almost all of the involvement with these compounds have been purely academic. Recently, new interest has been generated in the carbonyls as a means for the plating of the parent metal. This paper is intended to survey some of these deposition techniques and to review the literature on the chemistry of these compounds. We will deal largely with the bonding and structure of unsuhstituted metal carbonyls. The Discovery of Carbonyls
I n the Solvay process for the production of sodium bicarbonate, ammonia, and carbon dioxide vapors are passed successively into a saturated solution of sodium chloride. Previously, process engineers noticed that nickel valves in the system corroded very rapidly. Investigation of this phenomenon led to the discovery of trace amounts of carbon monoxide in the carbon dioxide and to Mond's classic b c o v e r y of nickel tetracarbonyl in 1891. Discovery of other carbonyls took place shortly thereafter, except for the last one, V(CO)6, which was not discovered until around 1960. The known carbonyls are listedin Table 1. Although there are no known binary palladium or platinum carbonyls, these metals do form carbonyl halides, e.g., Pt(C0)2 C12. It should be mentioned here that ((rr-C8H&, Ti(C0)2, and Cu(C0)CI are also known (I,$).
metal salt and a Group IA metal in an inert solvent under CO pressure (3000 psi) at 80°C (lOO°C for Cr). After steam distillation they report yields of 15%. Several other methods with higher yields (80-90%) are reported by E. W. Abel(4). This is not meant to imply that low temperature and pressure systems may be ruled out. King, Stokes, and Korenowski (6) report 16-20% yields of Mn~(CO)lofrom the reaction of (methylcyclopentadienyl) manganese tricarbonyl with a boiling solution of sodium in diglyme in a CO atmosphere. I n addition, Clark, Whiddon, and Serfass (6) report making 1-100 g quantities of C O ~ ( C Ofrom ) ~ the reaction of aqueous solutions of CoC12, KOH, and KCN with CO. Carbonyls may atso be prepared by displacement reactions with other carbonyls. Abel states (4) that "metathetical reactions of iron pentacarbonyl with molybdenum pentachloride and with tungsten hexachloride gave, respectively, molybdenum hexacarbonyl (28% yield) and tungsten hexacarbonyl (85% yield)." Bonding
Looking at the MO description of GO, one can visualize the structure.
Preporalion of the Carbonyls
Nickel, iron, and cobalt carbonyls can be prepared from the direct reaction of carbon monoxide and the metal. Other carbonyls can be prepared by the reduction of the metal from a salt in the presence of CO under high temperature and pressure. For example, Podall and Shapiro (3) report the preparation of Group VIB met$ hexacarbonvls by the reaction of the Grourr VIB
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The lone pair of carbon electrons are more strongly directed than the lone pair of oxygen for it is almost entirely a diagonal C(2S)-C(2Pz) hybrid. The Hantibonding orbitals are low enough in energy that they may accept electron density from filled nonbonding metal d orbitals (providing the d orbitals are of the correct symmetry for overlap). When this occurs, the
metal-ligand bond strength increases and the C-0 bond order decreases. (The metal-ligand u bond is increased due to the removal of the unfavorable charge separation induced by the donor a bond.) This change in the C-0 bond order shows up in the infrared region of the spectrum. Carbon monoxide itself has a C--0 stretch frequency a t about 2155 cm-' whereas terminal metal carbonyl groups have a C-0 stretch a t around 2000 cm.-' Such a correlation can be used to study r-bonding propertie: of substituent groups; however, it should he rsrnblished that the vibrations obrrrved in the infrared region reflect changes in r-bond strength and not in the c-bond strengths. By plotting the overlap integrals for the u and r bonds in carbon monoxide versus the internuclear distance, Kettle (7) has shown that the u bond is relatively insensitive to vibration and that r-electron effects are responsible for the spectrum (see Fig. 1) (7). The solid line in Figure 1 rgpresents the equilibrium C--O bond length (1.17 A). The
Figure 1. Plot of P, which is proportional to the internudear dirtonse versus the n-overlop integral IS.,) ond the r-overlap integral (ST).
two dotted vertical lines represent discursions of 0.05 A which is the order of niagnitude of change in bond length of the C - 4 bond during vibration (7). Proof of this inverse relationship between the metalligand and C-0 bond order is demonstrated by the isoelectric and isostructural Ni(CO)r, [Co(CO)&, and [Fe(CO)a]2-. The only difference between the species here is the increased charge density as we go from nickel to iron. Table 2 indicates (4) this reciprocal relation. When the CO ligands are added to the metal the Sidgwick bonding scheme holds, i.e., the effective atomic number is equal to that of the next inert gas. This scheme holds true for all binary carbonyls except V(CO)6 if we assume that: terminal carbonyl groups each contribute a pair of electrons to the outer shell of the metal atom; bridging carbonyl groups each contribute one electron to each of the metal atoms; and covalent metal-metal bonds are formed by one electron Table 2.
from each metal atom (4). Let us now look a t the actual structure of the carbonyls. Structure
Taking the groups in order, the first binary carbonyl is the anomalous V(CO)6. Since it is one electron short of having a complete inert gas shell, it is considerably less stable than Cr(C0)6, Mo(C0)6, or W(CO)a, and is the only paramagnetic binary carbonyl. There are no references in the literature as to the structure of V(CO)6, but it would seem'reasonable that it is a regular act* hedron. From infrared and Raman spectra and electron diffraction, it is well established that the Group VIB hexacarbonyls are perfect octahedra. Whether the CO bonds to the metal by a simple a bond in a d2sp" complex is questionable. In a paper by Pauling (8), it is shown that the valence bond theory predicts six a bonds with a bond order of unity for CI(CO)~while the molecular orbital theory suggests that with r bonding, the bond order would increase to a maximum of 1.5. The molecular orbital model seems more likely because the observed bond length for $he Cr-C bond is 1.916 + 0.003 A which is about 0.10 A less than the calculated single bond length of 2.02 1. I n addition, if six u bonds were present, the chromium atom would have a -6 formal charge. Since the electronegativity rule (Langmuir 1921, Pauling 1948) states "that the electric charge on each atom in a stable structure differ from zero by no more than one unit," (8) even correcting for the partial ionic character does not reduce the formal charge significantly. As Pauling points out (8) The electraneg&vity difference of Cr (electronegativity 1.6) and C (electronegativity 2.5) corresponds to only 19% double-bond character, which would reduce the charge an the chromium atom from 6- to 5-; no reasonable amount of partial ionic character of the single bonds would lead to agreement with the electronegativity principle.
Pauling further points out that if each bond had 50% double bond character, all nine outer orbitals (d6spS) could be utilized to form the twenty resonance forms represented in Figure 2. Although the formal charge on the chromium is -3, a 22% ionic character would reduce the formal charge to -1. Notice that this percent ionic character is close to the value 19% which corresponds to the diierence in electronegativity of chromium and carbon (8). The structures of the Group VIIB carbonyls may be typically represented by the structure of Mnz(C0)lo
Bond Numbers for the Species Ni(CO)* [Co(CO),]; and [Fe(COhIZNc-o
Ni(CO), [CO(CO)~I IFeiCO)rl'-
2.64 2.14 1.85
Mu-c 1.33 1.89 2.16
ZN 3.97 4.03 4.01
Figure 2.
Representotion of one of t v e d y . resononce forms of C I ( C 0 ) r
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Figure 5. The structures of %lC0)1r and FsdC0)n (9)
decreased tendency of the second-and'thiid row transition elements toward pentacoordination (4, 9).) Upon irradiation of the liquid Fe(CO), with uv light, golden yellow platelets of insoluble Fez(CO)r precipitate. It is now accepted that the structure of Fe2(CO)g involves bridging carbonyls which reflect none of the properties of "normal ketonic" carbonyls (see Fig. 4). Although they were once thought to exist, the enneacarbonyls of ruthenium and osmium do not (4,9).
quantities of the diamagnetic Fe&O)& are formed (see Fig. 6). The unique feature here is the presence of a pentacoordinated carbon atom, equidistant from all five iron atoms, just below the plane of the base. Due to the presence of the carbon atom Braye (10) states, all five iron atoms attain a closed shell structure in the ground state. This is d i c u l t to rationalbe because if we assume each Fe atom to be a daspahybrid, for the apexal Fe, all available orbitals are used up without considering any contributions from the carbon atom.
i/
where
Figure 4. F e d C 0 ) ~(9).
f!,
The FealC0)a species is ~roducedbv a varietv of reactions &lvin;( the caibonylferraie anion "(~e(CO)4z-)
I :s: .-l II
+C=~e=c=ij
and hydriodocarbonyliron (HFe(CO)4-). It is believed that Osa(C0)iz and Rua(C0)lz are formed by similar reactions. The Osa(C0)12 and Rua(C0)u have similar Dan structures, but Fea(CO)lzappears to have bridging carbonyl groups (see Fig. 5). The bonding in the 0 s and Ru structures can be visualized if one considers three octahedra coming together. The Fe structure to date is still hypothetical (4,9). If Fea(C0)12reacts with certain acetylenes, small 374 / Journal of Chemical Education
repream elestmnr domtd fmm ths n&ghbwing Fe ataa
reprarenb p i n of elashon. from CO
The apexal Fe's eight valence electrons can be envisioned as being distributed over the two lower energy d orbitals and four higher energy suborbitals of the daspahybrid orbitals. (Four electrons are paired in the lower de orbital and four are unpaired in the claspS orbital.) The four neighboring Fe atoms can now each contribute awelectron to pair up the unpaired electrons. The remaining suborbitals can be filled by the donation of electron pairs from the three neighboring CO groups. Perhaps the MO treatment of Braye would be more applicable (4). The cobalt family also forms several carbonyls, usually in the forms Mz(C0)s and h&(CO)lz (see Fig. 7). The Mz(C0)s structure may be considered identical with Fe2(CO)owith the exception that one CO bridging group is missing. I n determining the structure of Co&0)1~ an important point was brought out. The high resolution infrared spectrum was much simpler than the structure indicates. As Abel states (4):
%Wre ? .
a -
n o dNIhlrsr d =, &OW
rats-I
and b, eo!dl eledaaa.
dfTi&pse~snteis &d&gm e Ot haw s fdwe to obmm aamany inftwedabsosptiom &S amin p&eiple far a psdtiaular mdd is n& wu&l*e &-
bend &wa&w and b-rBh*m
&ace apmtmeha rnadel.lf Shodsun &o om aYtr(OCL)uwhasawhwtw~k net xet ht and E&IC&, Echr@CQl~,which haa two t m m M CO g m p s atbched tp w h met& hm four t.iails bridging C20 groups over &teame f@.@ef the wahebrallr tmiwsb m a atom W 3%. 8 $4, @I. As mdbned b h , hhe u d b b q eearbenyl in the niekd p n p icinielre~ Arhapa ~1b due tise Mwmqein iod.zati0P po%s~Vhl@ (d'o -+ i@ Of the S& pairerl &at@mi94%; Pa &Sj Ph 8.8 efl. !Fhe &&&ion potentid B inlfatlve Qf @e $U.& &orbit&" tenBemy to b& bmd. X-ray, iafr&'redv eleaaran difiCt&kin ~c&rs all indie& m
