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This question already has an answer here:

My A-Level Chemistry textbook assures me that:

  • a sigma bond is formed when two s orbitals overlap.
  • a pi bond is formed when two p orbitals overlap.
  • all single bonds between carbon and any other atom are sigma bonds
  • (side note) double bonds between carbon atoms consist of a sigma bond plus a pi bond
  • (side-side note) triple bonds between carbon atoms consist of a sigma bond plus 2 pi bonds.

However, the electronic structure of carbon is [He]2s22p2, implying two s orbitals and two p orbitals. I deduce from this that e.g. methane (CH4) has 2 sigma bonds and 2 pi bonds.

I am pretty sure that my deduction is incorrect, as it contradicts the statement in bold above, but cannot see why.

(In addition, the book states that methane has a completely regular tetrahedral structure with all four C – H bonds exactly the same, at identical angles to each other.)

Edit: Although aspects of this question overlap(!) with Can a s orbital overlap with any p orbital to form a sigma bond?, I am most interested in why what appears to a mixture of s- and p- bonds ends up creating such a symmetrical molecule as CH4.

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marked as duplicate by Mithoron, A.K., Todd Minehardt, Tyberius, Ben Norris Apr 13 at 13:40

This question has been asked before and already has an answer. If those answers do not fully address your question, please ask a new question.

  • $\begingroup$ Sigma bonds can involve any type of orbitals, as long as they have right geometry. Those in methane can be considered to be made from sp3 from C and s from H. $\endgroup$ – Mithoron Apr 9 at 21:46
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    $\begingroup$ Now it is just about time that someone utters the word "hybridization". $\endgroup$ – Ivan Neretin Apr 9 at 22:15
  • $\begingroup$ You'll get your answer here satisfactorily: chemguide.co.uk/basicorg/bonding/methane.html $\endgroup$ – Kartik Apr 10 at 16:51
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    $\begingroup$ @Kartik Many thanks for the link. Lots of really clear explanations there - I spent a couple of fascinated hours reading the site. $\endgroup$ – marktwo Apr 11 at 11:38
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I am going to attempt to answer this, based on the link supplied by @Kartik and expanding on @Mitheron’s comment above.

As we know, covalent bonding is the process of “filling up” partially filled outer orbitals by sharing electrons between atoms. Carbon’s structure [He]2s22p2, or, more accurately, [He]2s22px12py1. When bonding occurs, however, two things happen:

  • 1. One of the 2s electrons gets “promoted” to a p orbital. This is an “excited” state in which the electrons have slightly more energy (though not enough to remove any of them). The configuration therefore becomes [He]2s12px12py12pz1, and voila, we now have 4 unpaired electrons available to bond with other atoms.
  • 2. The extra energy allows the orbitals to “hybridize”. In effect, the 2s and 2p orbitals morph into 4 identical orbitals of slightly higher average energy. These orbitals are called sp3 orbitals (one part s, three parts p). The electronic structure of the ready-to-bond atom is now [He] (sp3)4.

These four sp3 orbitals of course repel each other equally, and hence end up pointing towards the vertices of an imaginary tetrahedron. Hence the regular structure of CH4 and analogous compounds.

Returning to the question, the sigma bonds formed by carbon are neither from overlapping s or overlapping p orbitals, but from overlapping sp3 orbitals.

The thing that makes a bond a sigma bond is not the type of subshell from which the electrons “originated”. It is the fact that the orbital overlap directly, with the two atomic nuclei and the main concentration of the shared electron cloud in a straight line. Pi bonds, on the other hand, overlap “sideways”, with the main concentrations of the bonding electron cloud off to the sides of the bonded atoms.

Please chime in, people, if I have omitted or misunderstood anything.

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