[Cre8ion]

Hi, I would gladly appreciate any help on the following questions:

1) Since each orbital holds 2 electrons, why do we only consider 1 electron to be shared per atom in a single covalent bond?

2) Why does an unfilled 3d orbital have more energy than that of a 4s orbital?

3) What are the factors affecting the strength of covalent bonds?

4) What is meant by an effective orbital overlap?

Thanks! 

[Hunter2]

1. Misunderstanding. The sharing of each orbitals electron gave 2 electron. A filled orbital don't do bonding. C-C . Each carbon gave 1 electron what gives 2 for the overlapping orbitals.
2. Its nature
3.-4 do your own attempt.

[Cre8ion]

1. Misunderstanding. The sharing of each orbitals electron gave 2 electron. A filled orbital don't do bonding. C-C . Each carbon gave 1 electron what gives 2 for the overlapping orbitals.

The electronic structure of an unbonded C atom 1s2 2s2 2p2, if it is a C-C single bond, that would be a sigma bond, between the 2s orbitals... unless it is between the 2p orbitals? 

And another question: When they state that electrons are added to an orbital in a particular manner (for e.g. a 4s orbital before 3d orbital), do they mean the adding of electrons across the period to (ie. from B to C the next electron is added to another 2p orbital), or does it also refer to the how electrons will be added to a particular element (during the formation of ionic bonds or covalent bonds)?

Thanks 

[Hunter2]

In bonding of a C it is hybridized. it will be sp³. All orbitals carrying 1 electron and share it with an orbital to the next atom, what means maximum two electrons in the overlapped orbitals.
Of course non bonding orbitals like in nitrogen and oxygen carrying 2 electrons from the same atom.

The orbitals are filled one by one first half filling then double filling. According Pauling and Hund law.  In some energetic cases the s orbital has to be filled first of the next shell, before the d and f-orbitals can be filled.

[magician4]

And another question: (...)

it refers to both: atoms (in the beginning of studies) are considered to be "hydrogen-like" with respect to their orbitals, and those are filled identically, no matter what the nucleus is

whether a given electronic configuration by the end of the day will result for example in an C^2- , a N^- or an oxgen-atom  (as the electronic configurations are identical)  is only a matter of what nucleus you consider to be at the core of that very electronic configuration

if you now, for example, take away one electron from each of the species mentioned above, you will result in  C^- , a nitrogen atom, and O⁺
...and again, the electronic configuration would be identical for all of them (but different from the one for C^2-, N^- and O  , of cause)

regards

Ingo

[Cre8ion]

@hunter2 @ingo thanks a lot! 

[Cre8ion]

A further question on the topic...

Since empty 3d orbitals are supposed to have more energy than empty 4s orbitals, when promoting a 3s orbital in the formation of covalent bonds (say in PCl5), should the promoted electrons go to 4s orbitals instead, since it is of lower energy than 3d orbitals?

Thanks once again

[magician4]

PCl₅ is a 10e-species ("hypervalent"), and hence requires the use of d- orbital(s) to form hybrid orbitals of the sp³d type  (tbp , VSEPR type AX₅)

as a result, the former s- and p orbitals will have disappeared, and only those 5 hybrid-orbitals and the remaining d orbitals will be available

the energy of those hybrid orbitals being lower than the resp. energy of the remaining d-orbitals, the hybrids become occupied


regards


Ingo