I am really confused on hybridization and how to do the orbital drawings. For example for BeF2: The electron configuration of F is (1s^2)(2s^2)(sp^5). In my book it says that there are two electrons in the ground state of Be in 1s and 2s orbital shells. For B3, 1 electron from 2s goes to the 2p shell. How do I know that and how do I know how many electrons go in each orbital? Also how do I distinguish a sp2 and an sp3 orbital shell. I know you have to promote electrons to different orbitals but how do you promote? Is there a certain system or procedure? Also how do I distinguish localized and delocalized bonding? How do I know when one exhibits a certain type. I'm really confused. Please help asap!! Thank you; thanks!!!

This is tough to do on a computer but here goes. First of all, I THINK you intended to write F as 1s2, 2s2, 2p5 and not sp5. Do I assume you have no trouble with BeF2. I think of it as Be= 1s2, 2s2 with 1 of the 2s electrons going to make ONE F 1s2, 2s2, 2p6 and the OTHER electron going to make ANOTHER F atom 1s2, 2s2, 2p6. That way we complete the octet for two F atoms and use the two available electrons from the one Be atom we have.
Now to B. B is 1s2, 2s2, 2p1. What happens is the promotion of one of the 2s2 to a p orbital so the final "hybridized" state is 1s2, 2s1, 2p2. First we call it hybridized because there is no difference now between the 2s and the 2p electrons; they have been rearranged, so to speak, to identical hybrid orbitals (somewhere in between a pure s and a pure p). Second, we call it sp2 by counting; we have 1 s electron and 2 p electrons so we call it simply sp2. Why didn't we promote more than 1? We have only 2 electrons in the 2s2 orbital. We promote one of them and NOT both of them. To my knowledge two are never promoted but I guess stranger things have happened. Furthermore, we know sp2 hybrids are trigonal planar so we can predict to an extent what many of the B compounds will be. Third, when we have sp3 hydbridization, that comes about because we have promoted a s electron from an element with TWO existing p electrons; for example, carbon.
C is 1s2, 2s2, 2p2. If we hybridize that, we promote one of the 2s electron to a 2p so it becomes 1s2, 2s1, 2p3 and we call it guess what? Counting we get 1 + 3 = 4 or sp3. Isn't that simple?
The energy for the promotion comes from the energy of formation of the final product. How do you know where to promote? If you have s electrons it is only logical to promote to the NEXT higher level which is a p ALTHOUGH there are some instances where it is promoted to much higher levels (Co(II) does that). Finally, you need to realize that much of this comes from experimentally determined behavior, I hope this helps but please post a follow up if anything is not a little clearer. As a parting note I should state here that BeF2 is NOT an ionic compound, as my answer above might suggest, and sp hybrids are formed abd those are linear compounds as sp would suggest.

Understanding hybridization and orbital drawings can be challenging, but I will do my best to explain the concepts to you.

First, let's clarify the electron configuration of fluorine (F). The correct electron configuration is 1s2, 2s2, 2p5, not sp5. The numbers in the superscripts represent the number of electrons in each orbital. So, for F, there are 2 electrons in the 1s orbital, 2 electrons in the 2s orbital, and 5 electrons in the 2p orbital.

Now let's focus on the orbital configuration of beryllium (Be) in the molecule BeF2. Beryllium has an electron configuration of 1s2, 2s2. In this case, both the 1s and 2s orbitals are completely filled with 2 electrons each.

To form BeF2, one electron from the 2s orbital of beryllium is promoted (or excited) to the 2p orbital. This promotion allows for the formation of two bonds with two fluorine atoms. So, the final electron configuration of Be in BeF2 is 1s2, 2s1, 2p1.

The concept of promoting electrons to different orbitals is driven by the formation of stable compounds. In general, electrons are promoted if it leads to the formation of more stable or energetically favorable compounds. The specific orbitals to which electrons are promoted depend on the specific situation, element, and compound being considered.

Now let's talk about hybridization and distinguishing between sp2 and sp3 orbitals. Hybridization occurs when atomic orbitals mix to form a new set of equivalent hybrid orbitals. The number and type of atomic orbitals that mix determine the hybridization and shape of the molecule.

For example, when one s and two p orbitals mix, they produce three new sp2 hybrid orbitals. This results in a trigonal planar molecular geometry. An example of a molecule with sp2 hybridization is boron trichloride (BCl3).

On the other hand, when one s and three p orbitals mix, they produce four new sp3 hybrid orbitals. This results in a tetrahedral molecular geometry. An example of a molecule with sp3 hybridization is methane (CH4).

To determine the hybridization of an atom in a molecule, you need to count the number of sigma bonds and lone pairs around that atom.

- If an atom has three sigma bonds and no lone pairs, it is sp2 hybridized.
- If an atom has four sigma bonds and no lone pairs, it is sp3 hybridized.

Localized and delocalized bonding have to do with the sharing of electrons between atoms. In localized bonding, electrons are shared between specific atoms in a molecule. This is commonly seen in covalent bonds, where electrons are shared between two atoms.

In delocalized bonding, electrons are not localized between specific atoms but are instead spread out or shared over multiple atoms or molecular orbitals. This is commonly seen in molecules with resonance structures or in molecular orbitals that extend over multiple atoms, such as in conjugated systems.

Determining whether a molecule exhibits localized or delocalized bonding requires looking at the electron distribution and molecular structure of the molecule. Experimental data and theoretical calculations can help in identifying the presence of specific bonding types.

It's essential to note that understanding hybridization, orbital drawings, and bonding types often requires practice and further study of molecular orbital theory and valence bond theory. You can refer to textbooks, online resources, or seek assistance from a teacher or tutor to gain a deeper understanding of these concepts.