Answer:
The orientation of the orbitals, the energy difference between them, their orbital sizes, and distance between the orbitals
Explanation:
1. Orientation
For example, the overlap between two p orbitals is roughly proportional to the cosine of the angle between them. An angle of 0° gives the best possible orbital overlap An angle of 90° means that the orbitals are orthogonal and have zero overlap. An angle of 180° means they are out of phase. There will be destructive overlap, and you will get an antibonding π* orbital.
2. Energy difference
If two orbitals have similar energies, they will have good overlap. The greater the energy difference, the poorer the orbital overlap. For example, a 3p orbital is higher in energy than a 2p. It can overlap efficiently with another 3p orbital, but not as well with a lower-energy 2p orbital.
3. Size difference
As the energy of an orbital increases, so does its size. A carbon 2p orbital cannot overlap efficiently with a chlorine 3p orbital that is twice its size.
4. Distance
The greater the distance between two orbitals, the poorer is their overlap.
The following factors govern the degree of overlap of atomic orbitals on different atoms.
1) symmetry of the orbitals
2) Energy of the orbitals
3) Orientation of the orbitals
The two overlapping orbitals must have the correct symmetry. Orbitals that do not have the appropriate symmetry can not overlap in a manner that leads to bond formation.
For orbitals to overlap, they must be comparable in energy. Orbitals that are far apart in energy can not overlap effectively leading to bond formation.
The lobes of the orbitals must point in the right direction (they must have the correct orientation) in order to ensure effective overlap leading to bond formation.
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