Delta bond

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Formation of a d bond by the overlap of two d orbitals Delta-bond-formation-2D.png
Formation of a δ bond by the overlap of two d orbitals
3D model of a boundary surface of a d bond in Mo2 Dimolybdenum-Mo2-delta-bond-Spartan-HF-3-21G-3D-side.png
3D model of a boundary surface of a δ bond in Mo2

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]

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 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.

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Pi bond Type of chemical bond

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Sigma bond

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Pi backbonding

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Quintuple bond

A quintuple bond in chemistry is an unusual type of chemical bond, first reported in 2005 for a dichromium compound. Single bonds, double bonds, and triple bonds are commonplace in chemistry. Quadruple bonds are rarer but are currently known only among the transition metals, especially for Cr, Mo, W, and Re, e.g. [Mo2Cl8]4− and [Re2Cl8]2−. In a quintuple bond, ten electrons participate in bonding between the two metal centers, allocated as σ2π4δ4.

Quadruple bond

A quadruple bond is a type of chemical bond between two atoms involving eight electrons. This bond is an extension of the more familiar types double bonds and triple bonds. Stable quadruple bonds are most common among the transition metals in the middle of the d-block, such as rhenium, tungsten, technetium, molybdenum and chromium. Typically the ligands that support quadruple bonds are π-donors, not π-acceptors.

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Phi bond

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A sextuple bond is a type of covalent bond involving 12 bonding electrons and in which the bond order is 6. The only known molecules with true sextuple bonds are the diatomic dimolybdenum (Mo2) and ditungsten (W2), which exist in the gaseous phase and have boiling points of 4,639 °C (8,382 °F) and 5,930 °C (10,710 °F). There is strong evidence to believe that there is no element with atomic number below about 100 that can form a bond with a greater order than 6 between its atoms, but the question of possibility of such a bond between two atoms of different elements remains open. Bonds between heteronuclear systems with two atoms of different elements may not necessarily have the same limit.

Molybdenum(II) acetate

Molybdenum(II) acetate is a coordination compound with the formula Mo2(O2CCH3)4. It is a yellow, diamagnetic, air-stable solid that is slightly soluble in organic solvents. Molybdenum(II) acetate is an iconic example of a compound with a metal-metal quadruple bond.

Potassium octachlorodirhenate(III) is an inorganic compound with the formula K2Re2Cl8. This dark blue salt is well known as an early example of a compound featuring quadruple bond between its metal centers. Although the compound has no practical value, its characterization was significant in opening a new field of research into complexes with quadruple bonds.

Metal-metal bond

In inorganic chemistry, metal-metal bonds describe attractive interactions between metal centers. The simplest examples are found in bimetallic complexes. Metal-metal bonds can be "supported", i.e. be accompanied by one or more bridging ligands, or "unsupported". They can also vary according to bond order. The topic of metal-metal bonding is usually discussed within the framework of coordination chemistry, but the topic is related to extended metallic bonding, which describes interactions between metals in extended solids such as bulk metals and metal subhalides.

References

  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.