In chemistry, delta bonds (δ bonds) are covalent chemical bonds, where four lobes of one involved atomic orbital overlap four lobes of other involved atomic orbital. This overlap leads to formation of a bonding molecular orbital with two nodal planes which contain internuclear axis and go through both atoms.[1][2][3][4]
The Greek letter δ in their name refers to d orbitals, as orbital symmetry of δ bond is same as that of usual (4-lobed) type of d orbital when seen down bond axis. This type of bonding is observed in atoms that have occupied d orbitals with low enough energy to participate in covalent bonding, for example, in organometallic species of transition metals. Some rhenium, molybdenum, technetium, and chromium compounds contain a quadruple bond, consisting of one σ bond, two π bonds and one δ bond.
The orbital symmetry of δ bonding orbital is different from that of a π antibonding orbital, which has one nodal plane containing internuclear axis and a second nodal plane perpendicular to this axis between atoms.
The δ notation was introduced by Robert Mulliken in 1931.[5][6] The first compound identified as having a δ bond was potassium octachlorodirhenate(III). In 1965, F. A. Cotton reported that there was δ-bonding as part of rhenium–rhenium quadruple bond in [Re2Cl8]2− ion.[7] Another example of a δ bond is proposed in cyclobutadieneiron tricarbonyl between an iron d orbital and four p orbitals of attached cyclobutadiene molecule.
References edit
- ^ Cotton, F. A.; Wilkinson, G. (1988). Advanced Inorganic Chemistry (5th ed.). John Wiley. p. 1087–1091. ISBN 0-471-84997-9.
- ^ Douglas, B.; McDaniel, D. H.; Alexander, J. J. (1983). Concepts and Models of Inorganic Chemistry (2nd ed.). Wiley. p. 137. ISBN 9780471895053.
- ^ Huheey, J. E. (1983). Inorganic Chemistry (3rd ed.). Harper and Row. p. 743–744. ISBN 9780060429874.
- ^ Miessler, G. L.; Tarr, D. A. (1998). Inorganic Chemistry (2nd ed.). Prentice-Hall. p. 123–124. ISBN 978-0138418915.
- ^ Jensen, William B. (2013). "The Origin of the Sigma, Pi, Delta Notation for Chemical Bonds". J. Chem. Educ. 90 (6): 802–803. Bibcode:2013JChEd..90..802J. doi:10.1021/ed200298h.
- ^ Mulliken, Robert S. (1931). "Bonding Power of Electrons and Theory of Valence". Chem. Rev. 9 (3): 347–388. doi:10.1021/cr60034a001.
- ^ Cotton, F. A. (1965). "Metal–Metal Bonding in [Re2X8]2− Ions and Other Metal Atom Clusters". Inorg. Chem. 4 (3): 334–336. doi:10.1021/ic50025a016.