Book

"4. Carbon molecule (C 2 ): The electronic configuration of carbon is 1s2 2s2 2p2. There are 12 electrons in C2. The electronic configuration of C2 molecule, therefore, is

The bond order of C2 is ½ (8 – 4) = 2 and C 2 should be diamagnetic. Diamagnetic C2 molecules have indeed been detected in vapour phase. It is important to note that double bond in C 2 consists of both pi bonds because of the presence of four electrons in two pi molecular orbitals. In most of the other molecules a double bond is made up of a sigma bond and a pi bond. In a similar fashion the bonding in N2 molecule can be discussed.

O2

5. Oxygen molecule (O 2 ): The electronic configuration of oxygen atom is 1s2 2s2 2p4. Each oxygen atom has 8 electrons, hence, in O 2 molecule there are 16 electrons. The electronic configuration of O 2 molecule, therefore, is

From the electronic configuration of O2 molecule it is clear that ten electrons are present in bonding molecular orbitals and six electrons are present in antibonding molecular orbitals. Its bond order, therefore, is Bond order = [Nb – Na] = [10 – 6] =2. So in oxygen molecule, atoms are held by a double bond. Moreover, it may be noted that it contains two unpaired electrons in π ∗2px and π ∗2py molecular orbitals, therefore, O 2 molecule should be Paramagnetic, a p r e d i c t i o n t h a t c o r r e s p o n d s t o experimental observation. In this way, the theory successfully explains the paramagnetic nature of oxygen. Similarly, the electronic configurations of other homonuclear diatomic molecules of the second row of the periodic table can be written. In Fig. 4.21 are given the molecular orbital occupancy and molecular properties for B2 through Ne2. The sequence of MOs and their electron population are shown. The bond energy, bond length, bond order, magnetic properties and valence electron configuration appear below the orbital diagrams. "

Question

1. Seeing it in atomic orbital perspective the molecules electronic config should be KK 2s2 2p4 ie 1 filled and 2 single filled orbitals. But according to the book there is no single filled molecular bond or anti-bond so the molecule is diamagnetic. Is the atomic bonding and molecular orbitals totally unrelated ?

2. When Pie bond is formed doesn't both atoms' orbitals needs to be acting i.e. Px of Atmos A acts with Px of Atom B ? There Px of 1 Atom and Py of other Atmos forms 1 molecular bond each ?

O2 Question

1. Why does O2 form 1 bond and 1 antibond ?

2. O2 form a double bond and complete the valance but the book shows 3 bonds but says double..

3. Explain "Moreover, it may be noted that it contains two unpaired electrons in π ∗2px and π ∗2py molecular orbitals"

Edited August 30Aug 30 by HbWhi5F

16 minutes ago, HbWhi5F said:

Book

"4. Carbon molecule (C 2 ): The electronic configuration of carbon is 1s2 2s2 2p2. There are 12 electrons in C2. The electronic configuration of C2 molecule, therefore, is

The bond order of C2 is ½ (8 – 4) = 2 and C 2 should be diamagnetic. Diamagnetic C2 molecules have indeed been detected in vapour phase. It is important to note that double bond in C 2 consists of both pi bonds because of the presence of four electrons in two pi molecular orbitals. In most of the other molecules a double bond is made up of a sigma bond and a pi bond. In a similar fashion the bonding in N2 molecule can be discussed.

O2

5. Oxygen molecule (O 2 ): The electronic configuration of oxygen atom is 1s2 2s2 2p4. Each oxygen atom has 8 electrons, hence, in O 2 molecule there are 16 electrons. The electronic configuration of O 2 molecule, therefore, is

From the electronic configuration of O2 molecule it is clear that ten electrons are present in bonding molecular orbitals and six electrons are present in antibonding molecular orbitals. Its bond order, therefore, is Bond order = [Nb – Na] = [10 – 6] =2. So in oxygen molecule, atoms are held by a double bond. Moreover, it may be noted that it contains two unpaired electrons in π ∗2px and π ∗2py molecular orbitals, therefore, O 2 molecule should be Paramagnetic, a p r e d i c t i o n t h a t c o r r e s p o n d s t o experimental observation. In this way, the theory successfully explains the paramagnetic nature of oxygen. Similarly, the electronic configurations of other homonuclear diatomic molecules of the second row of the periodic table can be written. In Fig. 4.21 are given the molecular orbital occupancy and molecular properties for B2 through Ne2. The sequence of MOs and their electron population are shown. The bond energy, bond length, bond order, magnetic properties and valence electron configuration appear below the orbital diagrams. "

Question

1. Seeing it in atomic orbital perspective the molecules electronic config should be KK 2s2 2p4 ie 1 filled and 2 single filled orbitals. But according to the book there is no single filled molecular bond or anti-bond so the molecule is diamagnetic. Is the atomic bonding and molecular orbitals totally unrelated ?

2. When Pie bond is formed doesn't both atoms' orbitals needs to be acting i.e. Px of Atmos A acts with Px of Atom B ? There Px of 1 Atom and Py of other Atmos forms 1 molecular bond each ?

O2 Question

1. Why does O2 form 1 bond and 1 antibond ?

2. O2 form a double bond and complete the valance but the book shows 3 bonds but says double..

3. Explain "Moreover, it may be noted that it contains two unpaired electrons in π ∗2px and π ∗2py molecular orbitals"

I always think it helps to see an energy level diagram for this. Here is diatomic oxygen:

You can see that 3 out of the 4 p electrons on each atom contribute to one σ bond and two π bonds. This would make a triple bond, BUT for the fact that there are still 2 p electrons left over. These go into the next lowest energy vacant orbital which is π*, antibonding. There are two of these of equal energy, so they can occupy both singly, as that will minimise their mutual electrostatic repulsion. So you end up with a double bond (because the π* cancels one of them out) and 2 unpaired electrons, which makes the molecule paramagnetic.

In the case of nitrogen those extra electrons are not present, so diatomic nitrogen does indeed have a triple bond - and is diamagnetic, having no unpaired electrons.

@exchemist

Please recommend resources to learn molecular orbitals

I get what you are saying, I am having problem visualizing and comparing molecular and atomic orbitals. I have some questions in main thread.

Edited September 1Sep 1 by HbWhi5F

10 minutes ago, HbWhi5F said:

@exchemist

Please recommend resources to learn molecular orbitals

I get what you are saying, I am having problem visualizing and comparing molecular and atomic orbitals. I have some questions in main thread.

To check you have understood (and to show you are not a robot), could you please describe to me which MOs on that energy level diagram would be populated with electrons in the case of C₂ and what bond order would result. I'll wait for your response before addressing the new thread you have posted.

As to resources, there are plenty of books. Your teacher should be able to recommend a suitable one for your level. For on-line resources I often find this source is helpful: https://chem.libretexts.org

Edited September 1Sep 1 by exchemist

In one of your earlier threads I asked if you understood orbitals and advised to to take a couple of steps back to full understaning.

I did not receive an answer to this.

Do you understand that s, p, d etc orbitals are just a convenient fiction (model) and do not describe what is actually going on in molecules ?

@studiot I think you meant atoms. Yes I understand s p d are 3d spaces where electron are like to be found (ie wave function).

14 minutes ago, HbWhi5F said:

@studiot I think you meant atoms. Yes I understand s p d are 3d spaces where electron are like to be found (ie wave function).

Actually I don't exactly mean atoms either.

And do you know why electrons are likely to be found there ?

@Stuart electron-electron repulsion ?

51 minutes ago, HbWhi5F said:

@Stuart electron-electron repulsion ?

See -You can hold a conversation, and hopefully get something out of it.

I am glad you know something about electron - electron repulsion, but that is not the reason why you are most likely to find electrons in these orbitals.

electron-electron repulsion is the reason why things are so complicated however.

I also note that you and exchemist have been having a conversation about bonding and anti bonding orbitals.

I can't remember if I have said it to you but I am always saying this.

A system, in this case of particles, attempts to minimise its energy on account of the Principle of Least Energy.

So if we write an accounting for all the energy of such a system and try to find the places where it is at a minimum we find that this gives us the shape of the orbitals in space.

It also gives solutions that are quantised - that is offers certain 'preferred' energy levels.

However - and this is where electron-electron repulsion comes in - We have only been able to solve this equation (known as the Schrodinger wave eqaution) for the case of a single electron - that is a hydrogen atom or Lithium ion or Helium ion.

To take account of multi electron atoms we use a special model, which is a modified hydrogen atom.

All the pictures you see actually refer to a hydrogen atom and sometimes they tell you this in the small print and then move on quickly.

We know it is a sufficiently good model because it works in general and its predictions match observatiosn quite closely.

So - bearing this in mind - would you like to look further into the how and why we turn this model of isolated atoms into a model of chemically bonded molecules ?

@studiot Yes I would like to know about molecular bonds

8 hours ago, exchemist said:

To check you have understood (and to show you are not a robot), could you please describe to me which MOs on that energy level diagram would be populated with electrons in the case of C₂ and what bond order would result. I'll wait for your response before addressing the new thread you have posted.

As to resources, there are plenty of books. Your teacher should be able to recommend a suitable one for your level. For on-line resources I often find this source is helpful: https://chem.libretexts.org

I didn't understand molecular bonds also I think c2 should form 1 sigma 1 sigma antibond 1 pie bonds and no antibonds but according to the mainthread (book) it froms 2 pie bonds .

2 hours ago, HbWhi5F said:

@studiot Yes I would like to know about molecular bonds

I didn't understand molecular bonds also I think c2 should form 1 sigma 1 sigma antibond 1 pie bonds and no antibonds but according to the mainthread (book) it froms 2 pie bonds .

What book are you using, just out of interest?

Yes I agree that, based on the diagram I provided, C₂ ought to have one σ bond and the equivalent of one π bond, though I think one electron would go into each of the two, like with O₂, which would give two half π bonds and would make the molecule paramagnetic.

Your source seems to say that does not happen and instead all 4 electrons go into the 2 π bonds, in which case it would be diamagnetic. That would imply that in C₂ the σ bond formed by the p(z) orbitals lies above the 2 π bonds in energy, and so is not filled. This could be accounted for by the effect of mixing, which is the subject of your other thread. (I must admit I either never knew, or have forgotten, this subtlety - I got my degree in 1976😀).

But yeah the mixing of the 2 pairs of σ and σ* orbitals (originating from the 2s and 2p(z) atomic orbitals) will lead to a greater energy separation of the 4 levels with the lowest bonding one becoming even more stable, but the second bonding one being a bit higher than before, which could shift it above the two π orbitals.

I'll pop a link to the Libretext page on orbital mixing in your other thread, so suggest taking a look at that.

Edited September 1Sep 1 by exchemist

1 hour ago, exchemist said:

What book are you using, just out of interest?

Further to this I also asked what are studying, without an answer.

What I also don't understand is why the bonding in C₂ is so important if as it appears you are still at the elementarty stage and this bonding has not yet been resolved (as of 2023) as far as I can tell.
As a species, C2 has definitely been measure by traditional chemical means (molecular weight), But there is also evidence of other bonding schemes given by spectroscopy.

Traditionally we have unravelled the molecular orbitals for ethane, ethylenes and acetylene by considering appropriate hybridisation schemes.
These are, in fact, often used a standard examples.

9 hours ago, studiot said:

Further to this I also asked what are studying, without an answer.

What I also don't understand is why the bonding in C₂ is so important if as it appears you are still at the elementarty stage and this bonding has not yet been resolved (as of 2023) as far as I can tell.
As a species, C2 has definitely been measure by traditional chemical means (molecular weight), But there is also evidence of other bonding schemes given by spectroscopy.

Traditionally we have unravelled the molecular orbitals for ethane, ethylenes and acetylene by considering appropriate hybridisation schemes.
These are, in fact, often used a standard examples.

That's interesting about C₂. Do you have a link to some information about its properties? It would be particularly interesting to see if it has been possible to determine if the ground state is a triplet or a singlet.

But from a teaching point of view I suspect the idea is not so much about C₂ as to teach the Aufbau filling of MOs and, importantly, the associated issues about what the order of energy levels actually is in different molecules, and why. This orbital mixing, and then the effect of relative penetration of σ orbitals and the effect of that across a period, was new to me but makes perfect sense.

P.S. Apologies, I've found there is a Wiki article on it: https://en.wikipedia.org/wiki/Diatomic_carbon , which does indeed confirm the ground state is a singlet with full occupancy of both π orbitals. Apparently it is a green gas that is a major constituent of carbon vapour, e.g. produced in electric arcs and even in blue hydrocarbon flames and is found in the interstellar medium.

Edited September 2Sep 2 by exchemist