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I am currently teaching myself AP Chemistry, so am at a very basic level still.

I am studying chemical bonds now, but am wondering why the orbitals don't make a difference to the way atoms bond.

To explain further: Oxygen has 6 valence electrons, and so when drawing the Lewis dot structure, you would draw it with 2 pairs of electrons and two single electrons. This seems consistent with the orbitals of an Oxygon atom, because it has 1 pair of electrons in the 3s orbital, another pair in one of the 3p orbitals, and 2 single electrons in each of the remaining 3p orbitals. This also makes sense to me in terms of and Oxygen atom forming a covalent bond with another Oxygen atom, since the two single electrons in each of the remaining 3p orbitals will "hang out" together, completing the octet rule for the valence shell of both atoms.

But when I looked up the Lewis dot structure for Carbon, I felt confused. It seems to be drawn with 4 single electrons . This feels inconsistent with what we know about the orbitals of a Carbon atom. Given that there are 4 valence electrons, with 1 pair in the 3s orbital, and 2 more single ones in each of two of the 3p orbitals, surely we should draw it as: 1 pair of electrons, 2 single electrons.

But we draw it as 4 single electrons. I know I am probably being really daft and missing something obvious, but please can someone help explain a bit more about why we draw it this way? And why the orbitals' pairs don't seem to matter when it comes to bonds?

Thanks

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2 Answers 2

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The only thing that matters for the Lewis structure is the number of valence electrons and possibly extra positive or negative charges that subtract/add electrons. Carbon has 4 valence electrons. That is why you draw it with 4 dots. You should also draw oxygen with 6 dots, adding the electrons one side at a time. Then you should consider the electrons on possible partners, for example a second oxygen, and try to complete the octets by sharing electrons and then merge them to bonds and lone pairs.

The concept of Lewis structures is older than orbitals and not strictly related. You will only confuse yourself if you try to derive Lewis structures from orbitals.

Many concepts in basic chemistry are empirically derived rules of thumb; these rules are not strict laws as laws from physics. You will find that many rules that are taught in basic chemistry have exceptions and break down at some point. So be careful when you read about them and try to check when they uphold and when they don't. This point is sadly often understated or completely missing in simple texts and can lead to tremendous confusion if the rules are taken at face value and when one tries to apply them without considering their limitations. You can also have concepts that appear to be contradicting each other. This is also often due to the fact that one concept was originally invented and applied to answer a specific and limited question and never intended to be transferred to other topics/properties.

This can make general chemistry frustrating and is the reason why I personally very much prefer physical and theoretical chemistry.

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If you have happened to learn how to write electronic configurations according the shells (not the orbital) ex: Oxygen(6): 2K 8L 6M


For Lewis Structures: we draw the valence shell electrons (here 6)

For Orbital Structures: we draw the orbitals of valence shell (here of 2s, 4p)

For Compounds, when we draw Lewis structure, we use these valence electrons only, but when drawing orbital structures we find the hybridization and then draw the orbitals.

comment for any doubt. Thanks.

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