1.3 Valence Bond Theory and Hybridization | Organic Chemistry

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For the F2 example did you mean their valence electrons are in the P orbital not S?

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[@lauraj2815]

I have a question. Towards the beginning of the video, in the H-F and F-F bonds, would it not be overlap between s (for H) and sp3 (for F), instead of p for F? I thought F would have sp3 hybridization, and thus it would be an sp3 hybridized orbital overlapping with the s orbital from H as opposed to a pure p orbital from F overlapping with the s from H. And then F-F would have overlap between two sp3 hybridized orbitals, instead of overlap of two p orbitals.

[@ChadsPrep]

Great question Laura! It turns out that only atoms that are making more than one bond experience hybridization. With F only making one bond, it simply used the atomic orbital where we find its unpaired electron, a 'p' orbital. So while you might look at the Lewis structure and see 4 electron domains and think sp3, F is not hybridized and is just using a 'p' orbital. The worst part of all of this is that when you get to molecular orbital theory, you'll learn that while the idea of hybrid orbitals has great utility in predicting bond angles and molecular geometry, it isn't an accurate reflection of reality in other ways (predicting energy levels of orbitals). So the theory does have some explanatory power, but molecular orbital theory has even greater explanatory power and is 'more' correct. So hybrid orbitals may not even really exist :( Hope this helps!

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[@fredbuls3298]

I'm not sure I understand overlapping orbitals. In methane, for example, each H shares a pair of electrons with the C. In one of those pairs, do the two electrons have identical wave functions (except for spin)?

[@ChadsPrep]

Hey Fred! First of all, keep in mind that Valence Bond Theory has its shortcomings...it's an incomplete theory but still has its utility. But it sets the stage for molecular orbital theory (the next lesson I'll be releasing later today). The main idea was that when two atoms bond together their electron clouds overlap to form a giant electron cloud that encompasses both atoms. Note my use of the word cloud instead of orbital just to give it a little different feel. In Molecular Orbital theory we'll find out the when two orbitals (which are depictions of wave functions) overlap they actually combine to form a new orbital (a new wavefunction). So we could say that the overlap of the Atomic Orbitals leads to the creation of a new Molecular orbital. But you'll see that even this explanation is too simplistic as the overlapping atomic orbitals will actually result in the creation of TWO molecular orbitals: one bonding and one antibonding. And if you get to molecules having more than two atoms it gets even more complicated, but we will largely ignore that in organic chemistry with one exception when we study conjugated systems (usually in the 2nd semester). Hope this helps!

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[@mayarise4242]

Valence Bond Theory Concept: Valence bond theory explains covalent bonding through the overlap of atomic orbitals. Electrons are shared in these overlapping regions. Atomic Orbitals: 1s Orbital: Spherical shape. 2p Orbital: Dumbbell shape with a node at the nucleus. Overlap Types: Sigma (σ) Overlap: Occurs along the internuclear axis (the line connecting the two nuclei). Pi (π) Overlap: Occurs side-to-side and involves p orbitals only. Examples of Sigma and Pi Overlap H2 Molecule: Sigma overlap of two 1s orbitals. HF Molecule: Sigma overlap of hydrogen’s 1s orbital with fluorine’s 2p orbital. F2 Molecule: Sigma overlap of two 2p orbitals. Pi Overlap: Only involves p orbitals and occurs above and below the internuclear axis. Hybridization Concept: Hybridization is the mixing of atomic orbitals to form new hybrid orbitals that can form stronger, more stable bonds. Determining Hybridization: Electron Domains: The number of atoms bonded to and lone pairs around an atom. Types: sp3 Hybridization: Four electron domains, bond angles of 109.5°. sp2 Hybridization: Three electron domains, bond angles of 120°. sp Hybridization: Two electron domains, bond angles of 180°. Examples of Hybridization Methane (CH4): Carbon is sp3 hybridized with four sigma bonds, bond angles of 109.5°. Ammonia (NH3): Nitrogen is sp3 hybridized with three sigma bonds and one lone pair, bond angles slightly less than 109.5° due to lone pair repulsion. Formaldehyde (H2CO): Carbon is sp2 hybridized with three sigma bonds and one pi bond, bond angles approximately 120°. Acetylene (C2H2): Carbons are sp hybridized with a triple bond (one sigma and two pi bonds), bond angles of 180°. Practical Applications Sigma Bonds: Formed by overlapping hybrid orbitals or atomic orbitals. Pi Bonds: Formed by the sideways overlap of unhybridized p orbitals. Lone Pairs: Reside in hybrid orbitals if the atom is hybridized.

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If the negative side of the orbital hits the positive side of the other orbital then they don't form a bond? What makes it positive or negative?

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