Question 1: When hybridization takes place, how do all the hybridized orbitals have the same energy? Here is an extract from my book: sp2 Hybridisation in C2H4: In the formation of ethene molecule, one of the sp2 hybrid orbitals of carbon atom overlaps axially with sp2 hybridised orbital of another carbon atom to form C–C sigma bond. While the other two sp2 hybrid orbitals of each carbon atom are used for making sp2–s sigma bond with two hydrogen atoms. The unhybridised orbital (2px or 2py) of one carbon atom overlaps sidewise with the similar orbital of the other carbon atom to form weak π bond, which consists of two equal electron clouds distributed above and below the plane of carbon and hydrogen atoms. The last time I checked, carbon was sp3 hybridized, and what is the "unhybridized orbital"? Why does the hybridization of carbon change with the formation of multiple bonds? Question 2: PCl5 has a trigonal bipyramidal structure, I do not understand how this is possible. The electronic configuration of the valence shell of P is 3s2 3p3. In its exited state, (I think) one electron from the 3s orbital will jump to the 4s orbital. So, totally, there are 2 s orbitals and 3 p orbitals taking part in bond formation with the 5 Cl atoms. The p orbitals will be bonded mutually perpendicularly to each other, where as the 2 s orbitals will be bonded in a position in which there is minimum repulsion from the other bonded pairs of electrons. But, the trigonal bipyramidal structure does not satisfy this condition. Why? Question 3: What is the difference between axial and equatorial bonds? Question 4: Can someone give me a formula to calculate the hybridization which will work for all the compounds? (I know 3 formulae, but each one has limitations) Thanks in advance.