[Seiryuu]

So I'm doing structural chemistry right now, and a few weeks ago we started on Lewis structures. We learned about how the atoms bond to each other with their valence electrons and form octets and that proper Lewis structures have minimized formal charges. What I don't understand are bonding rules. Now most times one electron from atom A is shared with one electron from atom B to form a single bond. But there are times where two atoms from atom A form a bond between it and some other atom. Why is this?

Example: Oxygen has six valence electrons. When it has a single bond and completes its octet, it has one single bond and six valence electrons. Why is it like this?

[subasurf]

Seiryuu, on 20 October 2011 - 12:40 AM, said:

When it has a single bond and completes its octet,

It doesn't. The oxygen atom has 6 valence electrons in it's outer shell. If it forms a SINGLE bond with another oxygen atom then it is not achieving the 8 electrons in it's valance shell that is 'wants'. By forming a double bond with another oxygen atom, it is sharing two electrons from the other atom and this is what completes the octet. IF an oxygen atom started off with 7 valance electrons (like Fluorine for example) then a single bond would be sufficient to complete the octet. If you look at the Lewis structure for an F2 molecule you'll see this.

hope that helps!

[farmboy]

subasurf, on 22 October 2011 - 07:08 AM, said:

It doesn't. The oxygen atom has 6 valence electrons in it's outer shell. If it forms a SINGLE bond with another oxygen atom then it is not achieving the 8 electrons in it's valance shell that is 'wants'. By forming a double bond with another oxygen atom, it is sharing two electrons from the other atom and this is what completes the octet. IF an oxygen atom started off with 7 valance electrons (like Fluorine for example) then a single bond would be sufficient to complete the octet. If you look at the Lewis structure for an F2 molecule you'll see this.

hope that helps!

What you say is true, but I suspect the OP may be asking about dative covalent bonds (been years since I used that term lol) in which one of the bonding atoms 'provides' both bonding electrons. I think, though I'm far from certain, that this happens in instances where one atom is fairly electropositive and the other strongly electronegative meaning that both the bonding electrons are assosciated with the more electronegative of the two atoms.

[Seiryuu]

farmboy, on 22 October 2011 - 08:46 AM, said:

What you say is true, but I suspect the OP may be asking about dative covalent bonds (been years since I used that term lol) in which one of the bonding atoms 'provides' both bonding electrons. I think, though I'm far from certain, that this happens in instances where one atom is fairly electropositive and the other strongly electronegative meaning that both the bonding electrons are assosciated with the more electronegative of the two atoms.

This may be the case, though I'll have to recheck my Lewis structures and make sure that there isn't an overall charge of the compound.

[Chap]

Are you asking why some elements form bonds which doesn't obey the octet rule? The reason behind this is dependent on the compound involved. For example, transition metal elements can form different numbers of bonds as they have a partially filled 'd' orbital, which can get involved in the bonding. Remember that not all the Lewis structures obey the Octet rule.

examples : XeF6, PF5,SF4


Hope that clears things up. If it doesn't, please ask!

[hypervalent_iodine]

Chap, on 9 December 2011 - 05:21 PM, said:

Are you asking why some elements form bonds which doesn't obey the octet rule? The reason behind this is dependent on the compound involved. For example, transition metal elements can form different numbers of bonds as they have a partially filled 'd' orbital, which can get involved in the bonding. Remember that not all the Lewis structures obey the Octet rule.

examples : XeF6, PF5,SF4


Hope that clears things up. If it doesn't, please ask!

What you've said about transition metals are true, however the examples you give are not coordination complexes, they're hypervalent compounds. The concept of hypervalency was a matter of contention for a long time after the development of the Lewis-Langmuir theory of bonding, which considers the covalent bonding between two atoms as a 2-centre 2-electron model - i.e. it is concerned with two atoms (the centres), which share a bond consisting of two electrons. Hypervalent compounds such as the ones you've listed and many others used to be considered as merely the 'exception to the rule' following Lewis' seminal work and as a result, not much was done to study them until the 1940's and onwards.


In the 1950's, hypervalent compounds were found to adopt what is known as the 3 centre-4 electron bonding archetype. The 3c-4e model initially arose from a qualitative application of molecular orbital theory, which describes the linear combination of one p-orbital from a central atom and two ligand p-orbitals to generate three molecular orbitals. It looks like this:



(picture copied from my own research proposal)

In contrast to the localised two-electron bonds (i.e. a localised sigma or pi bond) described by Lewis and Langmuir, in which there are two molecular orbitals (the antibonding and bonding molecular orbitals; or AMO and BMO, respectively), a 3c-4e bond combines to give three molecular orbitals - a bonding molecular orbital, a non-bonding molecular orbital (NBMO; the HOMO), which contains a node at the central atom, and an antibonding orbital (the LUMO). This model has since been supported by various computational models and is currently a generally accepted archetype for hypervalent bonding.


Anyway, this isn't stuff that you ever need to worry about in undergraduate, since it isn't really taught outside of 'some compounds contain central atoms which have a valency exceeding 8'. Nevertheless, I find it interesting, so I thought I'd share.

[Chap]

hypervalent_iodine, on 10 December 2011 - 04:37 AM, said:

What you've said about transition metals are true, however the examples you give are not coordination complexes, they're hypervalent compounds. The concept of hypervalency was a matter of contention for a long time after the development of the Lewis-Langmuir theory of bonding, which considers the covalent bonding between two atoms as a 2-centre 2-electron model - i.e. it is concerned with two atoms (the centres), which share a bond consisting of two electrons. Hypervalent compounds such as the ones you've listed and many others used to be considered as merely the 'exception to the rule' following Lewis' seminal work and as a result, not much was done to study them until the 1940's and onwards.


In the 1950's, hypervalent compounds were found to adopt what is known as the 3 centre-4 electron bonding archetype. The 3c-4e model initially arose from a qualitative application of molecular orbital theory, which describes the linear combination of one p-orbital from a central atom and two ligand p-orbitals to generate three molecular orbitals. It looks like this:



(picture copied from my own research proposal)

In contrast to the localised two-electron bonds (i.e. a localised sigma or pi bond) described by Lewis and Langmuir, in which there are two molecular orbitals (the antibonding and bonding molecular orbitals; or AMO and BMO, respectively), a 3c-4e bond combines to give three molecular orbitals - a bonding molecular orbital, a non-bonding molecular orbital (NBMO; the HOMO), which contains a node at the central atom, and an antibonding orbital (the LUMO). This model has since been supported by various computational models and is currently a generally accepted archetype for hypervalent bonding.


Anyway, this isn't stuff that you ever need to worry about in undergraduate, since it isn't really taught outside of 'some compounds contain central atoms which have a valency exceeding 8'. Nevertheless, I find it interesting, so I thought I'd share.

I greatly appreciate your reply!!! Thank you very much for the information. Always happy to learn.