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.