Molecular orbital
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In chemistry, a molecular orbital is a region in which an electron may be found in a molecule.[1] MOs are introduced in qualitative and pictorial models of bonding in molecules, and specify the spatial distribution and energy of one (or a pair) of electrons. More precisely, they are found quantitatively as wave functions, mathematical solutions to the Schrödinger wave equation for a molecule, using an approximation known as the Hartree-Fock or Self-Consistent Field method.
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[edit] Overview
A molecular orbital specifies the spatial distribution and energy of one (or a pair) of electrons. Most commonly a MO is represented as a linear combination of atomic orbitals (the LCAO-MO method), especially in qualitative or very approximate usage. They are invaluable in providing a simple model of bonding in molecules.
Most methods in computational chemistry today start by calculating the MOs of the system. A molecular orbital describes the behavior of one electron in the electric field generated by the nuclei and some average distribution of the other electrons. In the case of two electrons occupying the same orbital, the Pauli principle demands that they have opposite spin. Necessarily this is an approximation, and highly accurate descriptions of the molecular electronic wave function do not have orbitals (see configuration interaction).
[edit] Qualitative discussion
For an imprecise, but qualitatively useful, discussion of the molecular structure, the molecular orbitals can be obtained from the "Linear combination of atomic orbitals molecular orbital method" ansatz. In this approach, the molecular orbitals are expressed as linear combinations of atomic orbitals.
Molecular orbitals were first introduced by Friedrich Hund and Robert S. Mulliken in 1927 and 1928.[2] [3] The linear combination of atomic orbitals approximation for molecular orbitals was introduced in 1929 by Sir John Lennard-Jones.[4]. His ground-breaking paper showed how to derive the electronic structure of the fluorine and oxygen molecules from quantum principles. This qualitative approach to molecular orbital theory is part of the start of modern quantum chemistry.
Some properties:
- The number of molecular orbitals is equal to the number the atomic orbitals included in the linear expansion,
- If the molecule has some symmetry, the degenerate atomic orbitals (with the same atomic energy) are grouped in linear combinations (called symmetry adapted atomic orbitals (SO)) which belong to the representation of the symmetry group, so the wave functions that describe the group are known as symmetry-adapted linear combinations (SALC).
- The number of molecular orbitals belonging to one group representation is equal to the number of symmetry adapted atomic orbitals belonging to this representation,
- Within a particular representation, the symmetry adapted atomic orbitals mix more if their atomic energy level are closer.
[edit] Examples
[edit] H2
As a simple example consider the hydrogen molecule, H2, with the two atoms labelled H' and H". The lowest-energy atomic orbitals, 1s' and 1s", do not transform according to the symmetries of the molecule. However, the following symmetry adapted atomic orbitals do:
| 1s' - 1s" | Antisymmetric combination: negated by reflection, unchanged by other operations |
|---|---|
| 1s' + 1s" | Symmetric combination: unchanged by all symmetry operations |
The symmetric combination (called a bonding orbital) is lower in energy than the basis orbitals, and the antisymmetric combination (called an antibonding orbital) is higher. Because the H2 molecule has two electrons, they can both go in the bonding orbital, making the system lower in energy (and hence more stable) than two free hydrogen atoms. This is called a covalent bond. The bond order is equal to the number of bonding electrons minus the number of antibonding electrons, divided by 2. In this example there are 2 electrons in the bonding orbital and none in the antibonding orbital; the bond order is 1, and there is a single bond between the two hydrogen atoms.
A Molecular Orbital Diagram of H2
[edit] He2
On the other hand, consider the hypothetical molecule of He2 with the atoms labelled He' and He". Again, the lowest-energy atomic orbitals, 1s' and 1s", do not transform according to the symmetries of the molecule, while the following symmetry adapted atomic orbitals do:
| 1s' - 1s" | Antisymmetric combination: negated by reflection, unchanged by other operations |
|---|---|
| 1s' + 1s" | Symmetric combination: unchanged by all symmetry operations |
Similar to the molecule H2, the symmetric combination (called a bonding orbital) is lower in energy than the basis orbitals, and the antisymmetric combination (called an antibonding orbital) is higher. However, in its neutral ground state, each Helium atom contains two electrons in its 1s orbital, combining for a total of four electrons. Two electrons fill the lower energy bonding orbital, while the remaining two fill the higher energy antibonding orbital. Thus, the resulting electron density around the molecule does not support the formation of a bond between the two atoms (called a sigma bond), and the molecule does therefore not exist. Another way of looking at it is that there are two bonding electrons and two antibonding electrons; therefore, the bond order is 0 and no bond exists.
A Molecular Orbital Diagram of He2
[edit] Noble gases
Considering a hypothetical molecule of He2, since the basis set of atomic orbitals is the same as in the case of H2, we find that both the bonding and antibonding orbitals are filled, so there is no energy advantage to the pair. HeH would have a slight energy advantage, but not as much as H2 + 2 He, so the molecule exists only a short while. In general, we find that atoms such as He that have completely full energy shells rarely bond with other atoms. Except for short-lived Van der Waals complexes, there are very few noble gas compounds known.
[edit] Ionic bonds
When the energy difference between the atomic orbitals of two atoms is quite large, one atom's orbitals contribute almost entirely to the bonding orbitals, and the other's almost entirely to the antibonding orbitals. Thus, the situation is effectively that some electrons have been transferred from one atom to the other. This is called a (mostly) ionic bond.
[edit] MO diagrams
For more complicated molecules, the wave mechanics approach loses utility in a qualitative understanding of bonding (although is still necessary for a quantitative approach). The qualitative approach of MO uses a molecular orbital diagram. In this type of diagram, the molecular orbitals are represented by horizontal lines; the higher a line, the higher the energy of the orbital, and degenerate orbitals are placed on the same level with a space between them. Then, the electrons to be placed in the molecular orbitals are slotted in one by one, keeping in mind the Pauli exclusion principle and Hund's rule of maximum multiplicity (only 2 electrons per orbital (opposite spins); have as many unpaired electrons on one energy level as possible before starting to pair them).
The hardest part is to construct the MO diagram. For a simple molecule such as H2, we draw the diagram like this:
-
- __ σ*
-
- __ σ
The σ indicates a sigma bonding orbital, while σ* indicates a sigma antibonding orbital. We know that the diagram looks like this because we know that two s orbitals will interact to form a σ bonding orbital and a σ antibonding orbital.
Now if one considers N2, one realizes that the two nitrogen atoms each have a filled 1s orbital, a filled 2s orbital, and three half-filled 2p orbitals. The 1s orbitals, being inner shell, do not interact (or, equivalently, they are not valence electrons, as explained by valence bond theory).
The two 2s orbitals do, however, interact to create a σs orbital and a σs* orbital:
-
- __ σs*
-
- __ σs
If we assume that the interatomic axis joining the two N atoms is the z axis, we find that the two 2pz orbitals are able to overlap lobe-to-lobe to create a sigma bond. The two 2px and two 2py orbitals, lying perpendicular to the z axis, interact to create four pi orbitals (two bonding, two antibonding).
Finally, we must decide the order of the orbitals. The 2s orbitals, since they were initially of lowest energy, interact to create the lowest-energy orbitals. The 2p sigma bonds must be stronger than the pi bonds, so we expect the σp orbital to be lower than the πp orbital. However, this is not the case, primarily because of hybridization mixing the 2s and 2p orbitals. However, we do have the expected order for the σp* and πp* orbitals:
-
- ___ σp*
- ___ ___ πp*
- ___ σp
- ___ ___ πp
-
- ___ σs*
- ___ σs
As promised, there are 8 orbitals, the sum of the number of atomic orbitals (4+4) which combined to create the molecular orbitals. The total number of electrons is then 10 (five valence electrons from each atom). Two go into the σs orbital; two go into the σs* orbital; four into the two πp orbitals, and two go into the σp orbital.
The sigma bond order is the total number of electrons in sigma bonding orbitals (4), minus the total number of electrons in sigma bonding orbitals (2), all over 2 giving (4-2)/2 = 1. There is the similar pi bond order, giving (4 - 0)/2. Adding these together gives the total bond order. In this case the lowest two orbitals "cancel out"; there is one sigma bond and two pi bonds. Dinitrogen therefore has a triple bond.
Finally, we know that diatomic nitrogen is diamagnetic since there are no unpaired electrons in the diagram.
Unfortunately, this diagram is not applicable to oxygen, fluorine, and neon. This is because due to the higher electronegativity of these elements, the formation of hybrid orbitals is less important and thus we get the "expected" order of energy levels:
-
- ___ σp*
- ___ ___ πp*
- ___ ___ πp
- ___ σp
-
- ___ σs*
- ___ σs
The observation that the formation of hybrid orbitals is much less energetically favourable for smaller, more electronegative atoms (which are found in the first row) is due to the energy difference in the atom between the 2s and 2p orbitals. This energy difference increases from left to right along a row and from top to bottom down a column of the periodic table, so is highest for fluorine, which has the lowest mixing of s and p in the MOs. Mixing is most important when the energy difference is small.
If we were working with diatomic oxygen, we would use this MO diagram. In this case, there would be 12 electrons to place into molecular orbitals; the first ten go into the five orbitals of lowest energy; the last two, however, occupy separate πp* orbitals. The bond order is reduced to 2 since this is an antibonding orbital; also, the two unpaired electrons make liquid oxygen paramagnetic, which is not explained by the localized electron model.
A further observation is that molecular orbital theory explains why the dicarbon molecule, C2, does not contain a quadruple bond in its ground state although it would complete the octet - there are four bonding orbitals, but the top three cannot be occupied before one antibonding orbital is occupied.
[edit] More quantitative approach
To obtain quantitative values for the molecular energy levels, one needs to have molecular orbitals which are such that the configuration interaction (CI) expansion converges fast towards the full CI limit. The most common method to obtain such functions is the Hartree-Fock method which expresses the molecular orbitals as eigenfunctions of the Fock operator. One usually solves this problem by expanding the molecular orbitals as linear combinations of gaussian functions centered on the atomic nuclei (see linear combination of atomic orbitals and basis set (chemistry)). The equation for the coefficients of these linear combinations is a generalized eigenvalue equation known as the Roothaan equations which are in fact a particular representation of the Hartree-Fock equation.
Simple accounts often suggest that experimental molecular orbital energies can be obtained by the methods of ultra-violet photoelectron spectroscopy for valence orbitals and X-ray photoelectron spectroscopy for core orbitals. This however is incorrect as these experiments measure the ionization energy, the difference in energy between the molecule and one of the ions resulting from the removal of one electron. Ionization energies are linked approximately to orbital energies by Koopmans' theorem. While the agreement between these two values can be close for some molecules, it can be very poor in other cases.