1. Structure & Properties of Matter
Lesson # 6: Hybrid Orbitals

2. Orbitals & Bonding
We have already learned about determining where electrons could be in an atom based on probability and energy levels.
Knowing n and l, we can roughly determine the energy level and shape of the orbital.
It gets more complicated in trying to know what orbitals would look like in a bond between two atoms, because it would result from two orbitals overlapping as they share a pair of electrons.

3. Valence Bond Theory
Valence bond theory states that atomic orbitals overlap in bonding to form a new orbital with a pair of opposite spin electrons.
Valence bond theory is the second wave mechanical based theories of chemical bonding, along with the molecular orbital theory.

4. Example – Hydrogen
The simplest orbital overlap is in hydrogen (H2). Each has one electron in a 1s orbital. When the two atoms bond, the electron orbitals overlap, and the individual sphere shapes become closer to an oval.
The orbital overlap is called a sigma bond (σ) – the electron density is between the nuclei of the atoms.
1s1 1s s2
This works because each hydrogen atom has an empty spot for another electron in each of their 1s orbital. Any unfilled orbital can overlap with another unfilled orbital.

5. Example – HCl Hydrogen’s electron configuration is 1s1.
Chlorine’s valence electron configuration is 3s23p5. There is one spot in hydrogen’s s orbital to overlap with one of chlorine’s p orbitals. This is called an sp orbital.
1s 3p Overlap of 1s and 3p

6. Hybrid Orbitals
A hybrid orbital forms from the combination of at least two different orbitals.
Hybrid orbitals help explain resonance structures, exceeding the octet rule, and Lewis diagrams in general.

7. Example – Methane
We know from Lewis diagrams that carbon can make 4 bonds – one to each hydrogen in methane, as it has 4 electrons in its outer shell available for bonding.
BUT its electron configuration is:
AND its energy level diagram is: 2p 2s 1s
So wouldn’t it’s Lewis diagram look like this?

8. Methane (continued) 2p 2s 1s We know it actually looks like this:
This is due to orbital hybridization
One of the 2s electrons is promoted to the empty 2pz orbital, putting it in an excited state. This now allows carbon to make 4 bonds to hydrogen and make methane instead of just two. 2p 2s 1s

9. Methane (continued) 2sp3 1s
Doing this makes the energy of the 2s and the 2p the same, and is called a sp3 hybrid.
2sp3 1s

10. Example: Boron Trihydride2s 1s
We know Boron can make 3 bonds.
One of the 2s electrons must have been promoted to p.
This is called an sp2 hybrid.

11. Predicting Hybrid Orbitals based on VSEPR
Linear – sp
Trigonal Planar – sp2
Tetrahedral – sp3 (could also be trigonal pyramidal or bent if there are 4 orbitals used in bonding)
Trigonal bipyramidal – sp3d
Octahedral – sp3d2
(There are many more, but these are the ones we are focusing on in this unit.)

12. Double and Triple Bonds
Hybrid orbitals resulting from single bonds are called sigma bonds - σ
Hybrid orbitals resulting from double and triple bonds are called pi bonds – π
Pi bonds result from the electron density being concentrated in two separate regions that one on opposite sides of the two nuclei. The formation of pi bonds allows two atoms to share more than one pair of electrons between them.
VIDEO

13. Example - Ethene
The shape around each carbon is trigonal planar, meaning it can only do sp2 hybridization, which only accounts for 3 bonds. The two bonds to hydrogen, and one carbon-carbon bond make up the sp2, and they are sigma bonds.

14. Ethene (continued)
The fourth bond, the second carbon-carbon bond is a pi bond, and results from a p orbital overlapping either above or below the center on each carbon.
In summary, ethene has three hybridized sp2 sigma bonds, and one unhybridized 2p orbital – the pi bond.

15. Example - Ethyne
The shape around each carbon is linear, meaning it can only do sp hybridization, which only accounts for 2 bonds. The bond to hydrogen and one carbon-carbon bond make up the sp.
The third and fourth bonds, both carbon to carbon are pi bonds, and result from 2p orbitals overlapping from each carbon.
VIDEO

16. Summary
A covalent bond forms when an atomic orbital of one atom overlaps with an atomic orbital of another and a pair of electrons (with opposite spins) are shared between the overlapping orbitals.
Hybrid atomic orbitals form by mixing s, p, d orbitals with an atom before a bond occurs.
Bonds formed by hybrid orbitals are stronger than those formed by ordinary atomic orbitals.

17. Summary (continued)
Sigma bonds are formed by s-s, s-p, and end to end p-p, and even the overlap of two already hybridized orbitals.
Sigma bonds allow for free rotation around the bond axis since all the electron density is on the central atom (along the internuclear axis).
Pi bonds are formed by side to side overlap of p orbitals.
Pi bonds do not allow for free rotation as this would break or destroy the bond, as the electron density is above and below the internuclear axis.
A double bond is 1 sigma and 1 pi bond.
A triple bond is 1 sigma and 2 pi bonds.
VIDEO – Crash Course!