AO-AO Interactions

Atomic orbitals, solutions derived from the Schrödinger wave equation, are wavefunctions, where waves are well understood mathematically. Waves can be added together or subtracted from each other. Consider the addition of two sine waves. The product is a superposition, here:

Likewise, atomic orbitals can be added together to give a superposition called a molecular orbital or MO. Molecular orbitals are bonding when the orbital phase considerations are favourable:

The bonding MO wave function, Ψ (psi), can be squared, |Ψ|², to represent electron density. A bonding MO shows a build up of electron density between the two positively charged nuclei can be seen.

• The two positive nuclei are both attracted to the region of electron density and are shielded from each other.
• This, in essence, is the covalent bond.

If the atomic orbitals are "out of phase", the AO/AO interaction will exhibit a phase node – a regions in space with zero electron density – between the positive nuclei. There will be nothing to attract the nuclei together and the MO is (said to be) antibonding:

The diagram below shows the atomic orbitals, AOs, on a pair of adjacent atoms interacting with each other. Note that by convention we start reading from the bottom of the diagram because this is how MO diagrams are constructed, from the lower energies... up.

• First, we have the 1s + 1s -> sigma bonding MO, and then above this the sigma star antibonding MO.
• Then, there are the pair 2s + 2s bonding (sigma) and antibonding (sigma star) MOs.
• Next we come to the p + p interactions. These are a little more complicated because p orbitals are orientated in space along the x, y and z dimensions.
• The 2px + 2px orbital interaction gives a sigma bonding MO.
• The 2py + 2py and 2pz + 2pz interactions are equivalent, they only differ in spatial orientation. Both give rise to pi bonding MOs.
• Next come a pair of pi star -antibonding MOs.
• Finally, we have the p2x + 2px sigma star antibonding MO.

The diagram of the bonding and antibonding MOs is shown below:

Electrons are added to the MOs (bonding and antibonding) using the same rules that are used for adding electrons to atomic orbitals:

• The aufbau principle, lowest energy MOs fill first, or fill the orbitals from below, IF there are large orbital energy differences.
• The Pauli exclusion principle, a maximum of two electrons per orbitals and these must be of opposite spin
• Hund's rule, when there are equal energy or "degenerate" orbitals, these fill one electron at a time before pairing begins. Or, if two orbitals are not very different in energy, near-degenerate, or even exactly degenerate, then two electrons will fill both orbitals with parallel spin.

Leads to:

And so:

Simple, except that when electrons are added, sometimes the relative energies of the orbitals can change...

Dihydrogen

We can now examine diatomic species. We will start with the interaction of two hydrogen atoms, each with a single electron in a 1s AO. The two atoms come together and the two electrons go into the sigma 1s MO with is bonding. H2 is known to exist.

For dihydrogen, H2, we can identify the frontier molecular orbitals (FMOs). The highest occupied molecular orbital (or HOMO) is the sigma 1s MO. The lowest unoccupied MO (LUMO) is the sigma star 1s MO which is antibonding.

Bond order is defined as the number of electrons in bonding MOs (for H2 this is two) minus the number of electrons in antibonding MOs (zero) divided by two. Thus, hydrogen has a bond order of 1.

Dihelium

Now consider two helium atoms approaching. Two electrons go into the sigma 1s bonding MO, and the next two into the sigma star antibonding MO.

As antibonding MOs are more antibonding than bonding MOs are bonding, He2 (dihelium), is not expected to exist

And dihelium has a bond order of zero:

Dilithium

Now consider two lithium atoms interacting. Dilithium, Li2, is known in the gas phase, it has a bond order of one, and it has HOMO+LUMO FMOs:

Diberyllium

Now consider two beryllium atoms interacting. Diberyllium, Be2, has a bond order of zero and is unknown:

Diboron

Now consider two boron atoms interacting where each boron has five electrons. The ninth electron will go into the pi 2py bonding MO and the tenth electron into the degenerate (equal energy) pi 2pz bonding MO.

Analysis shows that diboron, B2, will be a diradical:

with a bond order of 1. Diboron, B2, is known in the gas phase and it is paramagnetic.

Dicarbon

Now consider two carbon atoms interacting. Dicarbon, C2, will have full, bonding pi 2py and pi 2pz MOs.

Dicarbon has (two equal energy) HOMOs and a LUMO.

Dicarbon has a bond order of two and it is known in the gas phase.

The orbital shapes can be calculated using MO software, such as Spartan by Wavefunction:

Calculated at the HF 6-31G* level.

Note that dicarbon is VERY reactive and it will polymerise (depending upon the conditions) to diamond, graphite, C60 "buckyballs" and similar species, single walled carbon nanotubes (SWNT) and multiwalled carbon nanotubes

Dinitrogen

Now consider two nitrogen atoms interacting:

Dinitrogen, N2, is well known, it has a bond order of three, and it has HOMO+LUMO FMOs:

Dioxygen

Before we can consider two oxygen atoms reacting to give dioxygen, O2, it should be noted that there is a slight reordering of the MO energy levels so that the sigma 2px bonding MO is lower in energy than the pi 2py and pi 2pz MOs. And, correspondingly, the sigma star antibonding 2px MO rises in energy above the pi star 2py and pi star 2pz MOs.

The result is that dioxygen, O2, is a diradical.

Oxygen is a well known, stable, diradical, bond order 2.

The diradical nature of triplet oxygen causes it to be paramagnetic, here.

Difluorine

Fluorine dimerises to F2, a species with a HOMO and a LUMO, bond order 1.

Dineon

Neon, like helium, does not dimerise because the dimer is bond order zero.

Summary

The data can be summarised:

Bond order zero species, are all unknown, while bond order > 0 are all known... although not necessarily as room temperature stable diatomic materials. Lithium, boron and carbon diatomic 'polymerise' to metals or network covalent materials.

Heteronuclear Diatomic Molecules & Molecular Ions

MO theory can be used to describe heteronuclear diatomic molecules & molecular Ions such as:

• LiH     Lithium hydride
• HF      Hydrogen fluoride
• CO     Carbon monoxide
• CN^–   Cyanide ion
• NO⁺   Nitrosonium ion

As the electronegativity differences increases the interacting orbitals will be at different energies. The result is that the covalent bonding energy decreases, but this is counteracted by an increasing electrostatic +/– attraction that is not represented on MO diagrams (see the Klopman equation, elsewhere in this web book.)

The MO diagram for the diatomic carbon monoxide, CO, shows it to be isoelectronic with nitrogen, N2:

The heteronuclear diatomic ions cyanide ion, CN^–, and nitrosonium ion, NO⁺, are also electronic with nitrogen, N2, and carbon monoxide. The only difference between the MO diagrams are the relative energies of the orbitals.

Please note that the primary aim of this page has been produced to introduce the idea of molecular orbitals and to show that the stable homonuclear diatomic species possess either LUMO and HOMO frontier molecular orbitals, or are diradical species.

The following is a message from Dr Eugen Schwarz in Germany – for which I am most grateful – concerns the above analysis:

For more information, go here and/or here.

© Mark R. Leach 1999-2009

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