Delta_bond Knowpia

In chemistry, delta bonds (δ bonds) are covalent chemical bonds, where four lobes of one involved atomic orbital overlap four lobes of the other involved atomic orbital. This overlap leads to the formation of a bonding molecular orbital with two nodal planes which contain the internuclear axis and go through both atoms.^[1]^[2]^[3]^[4]

Formation of a δ bond by the overlap of two d orbitals

3D model of a boundary surface of a δ bond in Mo₂

The Greek letter δ in their name refers to d orbitals, since the orbital symmetry of the δ bond is the same as that of the usual (4-lobed) type of d orbital when seen down the 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 the δ bonding orbital is different from that of a π antibonding orbital, which has one nodal plane containing the internuclear axis and a second nodal plane perpendicular to this axis between the 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 the rhenium–rhenium quadruple bond in the [Re₂Cl₈]²⁻ ion.^[7] Another interesting example of a δ bond is proposed in cyclobutadieneiron tricarbonyl between an iron d orbital and the four p orbitals of the 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.

[1] Cotton, F. A.; Wilkinson, G. (1988). Advanced Inorganic Chemistry (5th ed.). John Wiley. p. 1087–1091. ISBN 0-471-84997-9.

[2] Douglas, B.; McDaniel, D. H.; Alexander, J. J. (1983). Concepts and Models of Inorganic Chemistry (2nd ed.). Wiley. p. 137. ISBN 9780471895053.

[3] Huheey, J. E. (1983). Inorganic Chemistry (3rd ed.). Harper and Row. p. 743–744. ISBN 9780060429874.

[4] Miessler, G. L.; Tarr, D. A. (1998). Inorganic Chemistry (2nd ed.). Prentice-Hall. p. 123–124. ISBN 978-0138418915.

[5] 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.

[6] Mulliken, Robert S. (1931). "Bonding Power of Electrons and Theory of Valence". Chem. Rev. 9 (3): 347–388. doi:10.1021/cr60034a001.

[7] 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.

[1]

[2]

[3]

[4]

[5]

[6]

[7]

Summary

See also edit

References edit