Question about sp3 hybridization

In summary: I think the main takeaway is that regardless of the model, both theories end up giving similar results for the bond in ethene. In summary, the need for sp2 hybridized carbon with only 3 hybridized orbitals and one unhybridized orbital is due to the formation of a π-bond between adjacent sp2 hybridized carbons. This bond cannot form if the carbons are constrained to be pyramidal and have sp3 orbitals. The energy of this bond is lower than the energy gained from rehybridizing the orbitals on the individual carbons. There is debate between the valence bond and molecular orbital theories on the exact nature of the double bond, but both give similar results.
  • #1
Crystal037
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Homework Statement
why can't a carbon be sp3 hybridized and still form an ethene
Relevant Equations
carbon in ground state has electronic configuration=1s^2 2s^2 2p^2
why is sp2 hybridized carbon is need with only 3 hybridized orbitals and one unhybridized orbital? The unhybridized orbital forms a bond anyway. Then why can't they be hybridized and form the same bond
 
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  • #2
What is the geometry of sp3 orbitals? What is the geometry of sp2 orbitals? What is the geometry of ethene?
 
  • #3
I know the geometry of sp3 orbitals is tetrahedral and sp2 is trigonal planar which is the geometry of ethene.
So is it purely based on experimental value than on p orbital shouldn't be hybridized to form a trigonal planar shape
 
  • #4
In adjacent sp2 hybridized carbons the unpaired electrons in the pz orbitals can pair to form a π-bond, lowering the energy of the system. If you constrain the ethene carbons to be pyramidal (so that the orbitals are sp3 hybridized), that bond can't form--or at least the overlap between the unpaired electrons will be smaller. At any rate, the bent structure will be higher in energy.
 
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  • #5
why the overlap will be smaller. I am just saying than instead of overlapping of the carbon's unhybridized orbital pz with that of another carbon. Can't the two carbon have their orbitals overlapped in the hybridized state? the energy of hybridized orbital will definitely be less than that of unhybridized pz orbital
 
  • #6
Ok I think maybe I understand what you're getting at, and it's a tricky distinction that's gotten a lot of attention in the research literature. If you're studying intro chemistry, it's probably best to simply accept orbital hybridization as empirical, but if you're still interested, we can do a deeper dive.

Pauling (the originator of the concept of hybridization) actually proposed that ethene's double bond was built from two sp3-hybridized CH2 units, sometimes called banana bonds, whereas Huckel was the originator of the notion that a double bond was a combination of a sigma and a pi orbital. Wikipedia has a decent overview of this debate:
https://en.wikipedia.org/wiki/Sigma-pi_and_equivalent-orbital_models
Long story short, from a computational point of view, both theories end up giving the same answer (which is really only an approximation to the actual answer).

One thing I should point out, though
Crystal037 said:
the energy of hybridized orbital will definitely be less than that of unhybridized pz orbital
This is more complicated than you might think. A carbon atom is a carbon atom, so adding up the energy for the 4 sp3 orbitals should give the same answer as adding up the energy for 3 sp2 orbitals and a p orbital. But also, you need to remember that you're dealing with the interaction of two carbon atoms to form a bond. So even if you have unpaired electrons in higher-energy pz orbitals, those orbitals interact to form a pi bond that is significantly lower in energy than the energy that you gain by rehybridizing the orbitals on the individual carbons.
 
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  • #7
One other thing: pedagogically, thinking about the double bond as a sigma-pi combination can help with wrapping your head around certain processes in organic chemistry. Alkyl radicals and carbocations generally have trigonal planar structures, so it's straightforward to assert that their singly occupied/free orbitals are unhybridized p orbitals, and it's easy to use this framework to justify stabilization of reactive radical/cation centers by neighboring groups.
 
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  • #9
DrDu said:
Strange article. It claims in the first sentence that it is about valence bond theory, but sigma - pi separation is a concept of MO theory.
Hm, I didn't notice that. It's actually a decent example of how MO and VB approach the same problem in different ways.
 

1. What is sp3 hybridization?

Sp3 hybridization is a type of hybridization that occurs when one s orbital and three p orbitals combine to form four sp3 hybrid orbitals. This type of hybridization is commonly seen in molecules with tetrahedral geometry, such as methane (CH4).

2. How does sp3 hybridization affect the shape of a molecule?

Sp3 hybridization results in a tetrahedral shape for molecules. This is because the four sp3 hybrid orbitals are arranged in a way that maximizes the distance between them, resulting in a 109.5 degree bond angle.

3. What is the difference between sp3 hybridization and other types of hybridization?

Sp3 hybridization is different from other types of hybridization, such as sp, sp2, and sp3d, because it involves the combination of one s orbital and three p orbitals. This results in four sp3 hybrid orbitals, while other types of hybridization result in different numbers of hybrid orbitals.

4. How does sp3 hybridization affect the reactivity of a molecule?

Sp3 hybridization can affect the reactivity of a molecule by changing the types of bonds that are formed. For example, molecules with sp3 hybridization tend to form single bonds, which are generally less reactive than double or triple bonds. Additionally, the shape of the molecule can also affect its reactivity.

5. Can sp3 hybridization occur in molecules with more than four atoms?

Yes, sp3 hybridization can occur in molecules with more than four atoms. This is because the number of sp3 hybrid orbitals formed is determined by the number of s and p orbitals available, not the number of atoms in the molecule. For example, sulfur hexafluoride (SF6) has six atoms but still undergoes sp3 hybridization to form its octahedral shape.

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