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Does hybridisation only occur with Carbon?

  1. Jan 9, 2017 #1
    I've been studying how hybridisation works, and in very single example I've come across, there is always Carbon as the atom that undergoes hybridisation. But this can't be right, can it?

    I mean, Nitrogen can form triple bonds, Oxygen can form double bonds, etc.
    I'd assume then that Nitrogen gets sp hybridised and Oxygen gets sp2 hybridised?

    Is this correct?
     
  2. jcsd
  3. Jan 9, 2017 #2

    TeethWhitener

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    Hybridization can occur with most atoms. For instance, the nitrogen in NH3 is sp3 hybridized. The reason we tend to focus on carbon is because hybridization plays such an important role in understanding bonding in organic chemistry.
     
  4. Jan 10, 2017 #3

    DrClaude

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    Hybridization is not related to the number of bonds, but the spatial geometry of those bonds.
     
  5. Jan 11, 2017 #4

    DrDu

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    The only place where hybridization occurs is in our minds. You can describe nitrogen compounds with or without hybridized orbitals on N. The results are mostly very similar. It is also possible to describe bonding in ethylene using so called banana bonds made from sp3 hybrids on C.
    For oxygen, hybridization becomes quite unfavourable due to the huge energy difference between 2s and 2p orbitals.
     
  6. Jan 11, 2017 #5

    James Pelezo

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    Hybridization is a term that surfaces in explaining the transitions that occur at the valence level during molecular bonding giving 'Molecular Compounds'. This event can theoretically be applied to any nonmetal element, but is most frequently applied to Carbon, Nitrogen, Oxygen and the Halogens. However, the principles can be applied to Silicon, Phosphorous, & Sulfur. The concept of hybridization can be found in most freshman level general chemistry textbooks in the chapters on Molecular Geometry under the Valence Bond Theory. Here's a post entered on 9/14/2016 on Hybridization, maybe this will help... https://www.physicsforums.com/threads/why-does-hybridization-occur.885346/#post-5570100 => Dr Pelezo 9/14/2016. Hope this helps.
     
  7. Jan 12, 2017 #6

    Borek

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    What stops it from being applied to metals?
     
  8. Jan 12, 2017 #7

    DrDu

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    Nothing, although I think that despite Paulings resonating VB theory of metals, VB theory is more appropriate for non-metals than for metals, which are better described in terms of MO theory. The reason is that in metals both ionisation energies and electron affinities are small, so that generating ionic contributions in covalent bonds is energetically neutral in metals, but disfavoured in non-metals. MO theory includes these ionic configurations automatically, while you need many resonance structures to include them in VB. In covalent non-metal compounds, the ionic structures are suppressed whence VB gives often a better description.
     
  9. Jan 12, 2017 #8

    James Pelezo

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    Actually, nothing ... Be and B, using the Valence Bond Theory, are hybridized to sp and sp2 orbitals to generate linear and trigonal planer geometries respectively to accommodate formation of compounds like BeX2 and BX3 where X = Halide. For Gp IA metals, showing hybridization would mean showing an s-orbital changing into a paramagnetic hybridized orbital which overlaps with a paramagnetic orbital of the non-metallic anionic fragment to generate the diamagnetic bonded pair. Ionization would then be based upon a highly polarized system that undergoes 'heterolytic bond clevage'. Such, in my humble opinion, seems to be over-kill for describing a simple process. There might be circumstances/studies where such would give a better understanding of the process. I guess it would be more accurate to say (in my post) that hybridization is 'most frequently' applied to nonmetallic elements in molecular compounds or polyions; but both the valence structure of the cation metal and anion nonmetal fragments could (and, maybe should) be shown to be hybridized to accommodate the bonding process being described.
     
  10. Jan 13, 2017 #9

    DrDu

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    Whatever a paramagnetic orbital may be...
     
  11. Jan 13, 2017 #10

    James Pelezo

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    Are you asking what is paramagnetic? If so, in an unbonded hybrid orbital, containing only 1 electron is said to be paramagnetic. When two paramagnetic orbitals overlap, a covalent bond is formed containing paired electrons (the more stable configuration). The pairing of the electrons for the purpose of forming a bond between two elements is referred to as diamagnetic. Diamagnetic stability (paired electrons) is greater than paramagnetic stability and is the most fundamental driving force in the process of chemical bonding.
     
  12. Jan 14, 2017 #11

    DrDu

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    At best, the electron is paramagnetic, but not the orbital.
     
  13. Jan 15, 2017 #12

    James Pelezo

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    An orbital is the character of the electron and the electron the character of the orbital. Such is defined by the electron's quantum state within the element. Such gives rise to the 4 quantum numbers that define the nature of a given electron and associated orbital character. In the ground state configuration the electrons can have 'known' orbital geometries defined by s, p, d, & f configurations. When bonding takes place the quantum states of the valence electrons transform into orbital configurations that accommodate electron pairing. The terms used to describe an orbital containing one electron is paramagnetic (which does mean the electron is paramagnetic), but after bonding;i.e., the pairing of the electrons forms a diamagnetic condition and the sigma bond, assuming symmetry exists about the nuclei undergoing bonding.

    Consider BeCl2 ...
    upload_2017-1-15_1-10-55.png
     
  14. Jan 16, 2017 #13

    DrDu

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    Do you have a reference for this terminology?
     
  15. Jan 16, 2017 #14

    James Pelezo

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    The following notes are from Ebbing & Gammon; General Chemistry, 9th edn., Houghton Mifflin Co, New York, Copyright 2009; pgs 309 - 311. Also see Glossary in same edition.
     
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