Pi bond: Difference between revisions
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The Greek letter '''π''' in their name refers to [[p orbital]]s, since the orbital symmetry of the pi bond is the same as that of the p orbital when seen down the bond axis. P orbitals usually engage in this sort of bonding. However, d orbitals can engage in pi bonding. |
The Greek letter '''π''' in their name refers to [[p orbital]]s, since the orbital symmetry of the pi bond is the same as that of the p orbital when seen down the bond axis. P orbitals usually engage in this sort of bonding. However, d orbitals can engage in pi bonding. |
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Pi bonds are usually weaker than [[sigma bond]]s because their orbitals go further from the positive charge of the [[atomic nucleus]], which requires more energy. From the perspective of quantum mechanics, this bond weakness is explained by significantly less overlap between the previously p-orbitals due to their parallel orientation. |
Pi bonds are usually weaker than [[sigma bond]]s because their orbitals go further from the positive charge of the [[atomic nucleus]], which requires more energy. From the perspective of quantum mechanics, this bond weakness is explained by significantly less overlap between the previously p-orbitals due to their parallel orientation. |
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Although the pi bond by itself is weaker than a [[sigma bond]], pi bonds are most often found in multiple bonds together with sigma bonds and the combination is stronger than either bond by itself. This can be seen from comparison of the carbon carbon [[bond length]]s in [[ethane]] with 154 [[picometer|pm]], [[ethylene]] with 133 pm and [[acetylene]] with 120 pm. |
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[[Image:Pi-bond.jpg|thumb|left|two p-orbitals forming a π-bond]] |
[[Image:Pi-bond.jpg|thumb|left|two p-orbitals forming a π-bond]] |
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Atoms with [[double bond]]s or [[triple bond]]s have one sigma bond and the rest are usually pi bonds. Pi bonds result from parallel orbital overlap: the two combined orbitals meet lengthwise and create more diffuse bonds than the sigma bonds. Electrons in pi bonds are sometimes referred to as pi electrons. |
Atoms with [[double bond]]s or [[triple bond]]s have one sigma bond and the rest are usually pi bonds. Pi bonds result from parallel orbital overlap: the two combined orbitals meet lengthwise and create more diffuse bonds than the sigma bonds. Electrons in pi bonds are sometimes referred to as pi electrons. |
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===Special cases=== |
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Pi bonds do not necessarily have to connect atoms; pi interactions between the metal atom and the sigma bond of molecular hydrogen play critical roles in the reduction of some organometallic compounds. [[Alkyne]] and [[alkene]] pi bonds often bond with metals in a bond that has significant pi character. |
Pi bonds do not necessarily have to connect atoms; pi interactions between the metal atom and the sigma bond of molecular hydrogen play critical roles in the reduction of some organometallic compounds. [[Alkyne]] and [[alkene]] pi bonds often bond with metals in a bond that has significant pi character. |
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Only in some molecules are the sigma bonds more energetic than the pi bonds. Examples are [[diironhexacarbonyl]] (Fe<sub>2</sub>(CO)<sub>6</sub>), [[dicarbon]] (C<sub>2</sub>) and the [[borane]] B<sub>2</sub>H<sub>2</sub>. In these compounds the central bond consists only of pi bonding and in order to achieve maximum orbital overlap the bond distances are much shorter than expected <ref>''Bond length and bond multiplicity: σ-bond prevents short π-bonds'' Eluvathingal D. Jemmis, Biswarup Pathak, R. Bruce King, Henry F. Schaefer III [[Chemical Communications]], '''2006''', 2164 - 2166 [http://dx.doi.org/10.1039/b602116f Abstract]</ref>. |
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== See also == |
== See also == |
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* [[Molecular geometry]] |
* [[Molecular geometry]] |
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* [[Sigma bond]] |
* [[Sigma bond]] |
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==References== |
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<references/> |
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[[Category:Chemical bonding]] |
[[Category:Chemical bonding]] |
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Revision as of 16:28, 28 May 2006
In chemistry, pi bonds (π bonds) are chemical bonds of the covalent type, where two lobes of one involved electron orbital overlap two lobes of the other involved electron orbital. Of the orbital's node planes, one only goes through both atoms.
The Greek letter π in their name refers to p orbitals, since the orbital symmetry of the pi bond is the same as that of the p orbital when seen down the bond axis. P orbitals usually engage in this sort of bonding. However, d orbitals can engage in pi bonding.
Pi bonds are usually weaker than sigma bonds because their orbitals go further from the positive charge of the atomic nucleus, which requires more energy. From the perspective of quantum mechanics, this bond weakness is explained by significantly less overlap between the previously p-orbitals due to their parallel orientation.
Although the pi bond by itself is weaker than a sigma bond, pi bonds are most often found in multiple bonds together with sigma bonds and the combination is stronger than either bond by itself. This can be seen from comparison of the carbon carbon bond lengths in ethane with 154 pm, ethylene with 133 pm and acetylene with 120 pm.
Atoms with double bonds or triple bonds have one sigma bond and the rest are usually pi bonds. Pi bonds result from parallel orbital overlap: the two combined orbitals meet lengthwise and create more diffuse bonds than the sigma bonds. Electrons in pi bonds are sometimes referred to as pi electrons.
Special cases
Pi bonds do not necessarily have to connect atoms; pi interactions between the metal atom and the sigma bond of molecular hydrogen play critical roles in the reduction of some organometallic compounds. Alkyne and alkene pi bonds often bond with metals in a bond that has significant pi character.
Only in some molecules are the sigma bonds more energetic than the pi bonds. Examples are diironhexacarbonyl (Fe2(CO)6), dicarbon (C2) and the borane B2H2. In these compounds the central bond consists only of pi bonding and in order to achieve maximum orbital overlap the bond distances are much shorter than expected [1].
See also
References
- ^ Bond length and bond multiplicity: σ-bond prevents short π-bonds Eluvathingal D. Jemmis, Biswarup Pathak, R. Bruce King, Henry F. Schaefer III Chemical Communications, 2006, 2164 - 2166 Abstract