Understanding molecular stability

[lemur]

This is probably a completely elementary chemistry question, but I'm trying to figure out why molecules like O2, CO2, CO, H2O are stable. With H2, it makes since to me that the two H atoms combine to form a full outer shell, and thus resemble Helium. However, O has 8 electrons so O2 must have 16, which is 2 less than a full outer shell, so why doesn't it receive two more electrons? I understand that ionization also makes a difference and that charge-neutral molecules wouldn't have a reason to bond further but how does that work in practice? E.g. does O2 break apart into, say, one O cation and one O anion and then each receives compatible pairing partners? How do you figure out the process that will occur when compounds are mixed and heated, etc.? It makes sense to me that heat breaks bonds and frees electrons, but how do you predict the details of how and what exactly will happen?

[John Cuthber]

The oxygen atoms each have 8 electrons.

Two are "stuck" in an inner shell and don't participate in any chemistry.

The 6 that are left are each 2 short of the 8 needed to form a closed shell- so they share and get 8 each (albeit that they only have a half- share of two of the 8).

Does that help?

[lemur]

The oxygen atoms each have 8 electrons.

Two are "stuck" in an inner shell and don't participate in any chemistry.

The 6 that are left are each 2 short of the 8 needed to form a closed shell- so they share and get 8 each (albeit that they only have a half- share of two of the 8).

Does that help?

Oh, right. That's what those diagrams with the dots show. So when the O2 gets enough energy to break into atoms, each ends up with one of the two shared electrons, which is one more than it needs to balance its protons' charge? And so you get two negatively charged O ions by heating up O2?

 

edit: so if O needs two and C needs 4, how does CO become stable?

Edited April 27, 2011 by lemur

[mississippichem]

Oh, right. That's what those diagrams with the dots show. So when the O2 gets enough energy to break into atoms, each ends up with one of the two shared electrons, which is one more than it needs to balance its protons' charge? And so you get two negatively charged O ions by heating up O2?

 

A decent guess. This used to confuse me my freshmen year. You actually end up with two uncharged O atoms [diradicals, two unparied electrons]. Each has 8 electrons, as the [ce] O_{2} [/ce] they came from had 16.

 

This is actually pretty hard to do though, unless you have substrate to react with the free O atoms. as an [ce] O_{2} [/ce] molecule would be more likely to just gain an electron from somewhere else to form a superoxide radical anion: [ce] O_{2} \cdot ^{-} [/ce], a paramagnetic, 17 electron species. Some have postulated that this is the first step in the combustion of organic molecules in air.

 

[ce] O_{2} [/ce] can have some counter intuitive chemistry as it exists as a diradical in the ground state [two unparied electrons]. Molecular oxygen would rather spin pair one of its radicals than split into two more diradicals.

[lemur]

A decent guess. This used to confuse me my freshmen year. You actually end up with two uncharged O atoms [diradicals, two unparied electrons]. Each has 8 electrons, as the [ce] O_{2} [/ce] they came from had 16.

That makes sense. So should I look at is as each atom lending two to the other in the O2?

 

This is actually pretty hard to do though, unless you have substrate to react with the free O atoms. as an [ce] O_{2} [/ce] molecule would be more likely to just gain an electron from somewhere else to form a superoxide radical anion: [ce] O_{2} \cdot ^{-} [/ce], a paramagnetic, 17 electron species. Some have postulated that this is the first step in the combustion of organic molecules in air.

That's interesting. Why does the O2 only gain one electron instead of two to fill the shell? Is that because it is uncharged as a molecule and is thus drawing the extra electron purely because its outer shell has openings?

 

 

[ce] O_{2} [/ce] can have some counter intuitive chemistry as it exists as a diradical in the ground state [two unparied electrons]. Molecular oxygen would rather spin pair one of its radicals than split into two more diradicals.

Why is it called "spin pairing?" Maybe I should just google that one along with some of your other terms. You explain things well, though, thanks.

[John Cuthber]

Unfortunately, CO is complicated.

A simple approach is that the carbon ends up nearer where it wants to be and the oxygen is OK. The real answer involves back bonding and molecular orbitals. That's a bit complicated for this sort of thread.

[mississippichem]

That makes sense. So should I look at is as each atom lending two to the other in the O2?

 

It's a bit difficult to deterministically say how many electrons are being shared, remember we're dealing electron wave/particles. But certain approximation techniques show that the two oxygen atoms share 4 electrons. The two unpaired electrons can almost never be shared because they are in anti-bonding orbitals. The sharing of those two electrons is rare because that requires a parity violation, a rule that comes out of group theory used in quantum chemistry.

 

That's interesting. Why does the O2 only gain one electron instead of two to fill the shell? Is that because it is uncharged as a molecule and is thus drawing the extra electron purely because its outer shell has openings?

 

If there is an available proton around, [ce] H^{+} [/ce], the superoxide ion can grab that proton and accept another electron to form the peroxide ion: [ce] HOO^{-} [/ce] which really contains the [ce] OO^{2-} [/ce] fragment.

 

Why is it called "spin pairing?" Maybe I should just google that one along with some of your other terms.

 

All fermions, this includes electrons, have spin +1/2 or -1/2 [the sign choice is arbitrary]. Every orbital can accommodate two electrons, and each electron must have a unique set of quantum numbers. So electrons are spin paired if they both reside in the same orbital and have opposite spins.

 

You explain things well, though, thanks.

 

Glad to know someone out there in cyberspace is listening to my rants.

 

Unfortunately, CO is complicated.

A simple approach is that the carbon ends up nearer where it wants to be and the oxygen is OK. The real answer involves back bonding and molecular orbitals. That's a bit complicated for this sort of thread.

 

John Cuthber:

 

Physicists say a diatomic molecule is a molecule with one too many atoms. Physical chemists say that an asymmetric diatomic is a molecule with one too few n-fold axes of rotation .

 

Lemur:

 

Unfortunately whenever you have a molecule that is very asymmetric and with orbital energies or effective nuclear charges that are very different, the picture gets pretty complicated. I would be doing you a disservice by trying to explain it without evoking the math of group theory and overlap integrals. Deriving the orbital configurations for these molecules is laborious even for the experienced. In real life, you can just let Waveunction Spartan software do the work for you.

Edited April 27, 2011 by mississippichem

[lemur]

If there is an available proton around, [ce] H^{+} [/ce], the superoxide ion can grab that proton and accept another electron to form the peroxide ion: [ce] HOO^{-} [/ce] which really contains the [ce] OO^{2-} [/ce] fragment.

That's because the Hydrogen atom's electron is attracted to the O2 because its outer shell isn't full? What is the actual force mechanism that causes that attraction? Is there some slight surplus charge for some reason, like because of the distance of the electrons from the nucleus or something like that?

 

Also, when you talk about free Hydrogen atoms, wouldn't those only be available when other H-containing molecules are breaking down within the reaction? Free hydrogen just floats away into space, right? off-topic: does that mean Earth has a limited supply of hydrogen that is gradually being converted into water? There are no natural mechanisms where hydrogen gets extracted from water, are there? There's another potential cause for rising sea-level.