Atoms in the row of the periodic table running from Li to F have electrons in 2s and 2p orbitals, and as all molecules of interest to organic chemists contain at least one such atom we now need to think about how 2s and 2p orbitals interact. We also need to introduce you to a useful piece of terminology that is used to describe the symmetry of molecular orbitals. We can do all of this by thinking about the bonding in another ubiquitous diatomic gas, N2. N atoms have electrons in 1s, 2s, and 2p orbitals, so we need to consider interactions between pairs of each of these orbitals in turn.

1s orbitals we have already dealt with. Combining 2s orbitals is essentially just the same; they form bonding and antibonding orbitals just as 1s orbitals do and with similar shapes too, but higher energies, because the 2s orbitals are higher in energy that the 1s orbitals. The 2s orbitals are also bigger than 1s orbitals, and because of their ‘onion skin’ form, the exact nature of the MOs they give rise to is more complex than those which come from 1s AOs, but you can represent them in sketches in just the same way:

The bonding orbitals formed from 1s–1s and 2s–2s interactions have another important feature in common: they are all cylindrically symmetrical. In other words, if you look at the molecular orbital end-on, you can rotate it around the axis between the two atoms by any amount and it looks identical. It has the symmetry of a cigar, a carrot, or a baseball bat. Bonding orbitals with cylindrical symmetry like this are known as σ (sigma) orbitals, and the bonds which result from putting two electrons into these orbitals are known as σ bonds. The single bond in the H2 molecule is therefore a σ bond. The antibonding orbitals which result from combining these AOs are also cylindrically symmetrical and are called σ* orbitals, with the * denoting their antibonding character. Now for the 2p orbitals. As described on p. 86 each atom has three mutually perpendicular 2p atomic orbitals. In a diatomic molecule, such as N₂, these 2p orbitals must combine in two different ways—one p orbital from each atom (shown in red here) can overlap end-on, but the other two p orbitals on each atom (shown in black) must combine side-on.

We’ll deal with the end-on overlap fi rst. This is what happens if we combine the two 2p orbitals out of phase: as with the 2s orbitals, we have a node between the atoms, and any electron in this MO would spend most of its time not between the nuclei—as you can guess, this is an antibonding orbital.

If we combine the orbitals in phase, this is what we get.

There is a nice rich area of electron density between the nuclei, and somewhat less outside, so overall filling this orbital with electrons would lead to an attraction between the atoms and a bond would result. Both of these MOs have cylindrical symmetry and are therefore designated σ and σ* orbitals, and a bond formed by filling the MO made from interacting two 2p orbitals end-on is called a σ bond.

● σ bonds can be made from s or p atomic orbitals, provided they form a cylindrically symmetrical molecular orbital.

Each atom presents its other two 2p orbitals for side-on overlap. This is what the antibonding MO formed by out-of-phase combination of two side-on p orbitals looks like:

These MOs do not have cylindrical symmetry—in fact you have to rotate them 180° about the axis between the nuclei before you get back something looking like what you started with but with opposite phase—and as a result the symmetry of these orbitals is given the symbol π: the bonding orbital is a π orbital and the antibonding orbital is a π* orbital. Bonds which are formed by filling π orbitals are called π bonds, and you’ll notice that because of the π symmetry the electron density in these bonds does not lie directly between the two nuclei but rather to either side of the line joining them. Since an atom has three mutually perpendicular 2p orbitals, two of which can interact side on in this way, there will exist a pair of degenerate (equal in energy) mutually perpendicular π orbitals and likewise a pair of degenerate mutually perpendicular π* orbitals. (The third p orbital interacts end-on, forming a σ orbital and a σ* orbital, of course). The two sorts of MOs arising from the combinations of the p orbitals are, however, not degenerate—more overlap is possible when the AOs overlap end-on than when they overlap side-on. As a result, the 2p–2p σ orbital is lower in energy than the 2p–2p π orbitals. We can now draw an energy level diagram to show the combinations of the 1s, 2s, and 2p AOs to form MOs, labelling each of the energy levels with σ, σ*, π, or π* as appropriate.

Now for the electrons. Each nitrogen atom contributes seven electrons to the molecule, so we have to fill this stack of orbitals with 14 electrons, starting at the bottom. The result is:

The σ and σ* MOs formed from interactions between the two 1s orbitals, and the two 2s orbitals are all filled: no overall bonding results because the filled bonding and antibonding orbitals cancel each other out. All the bonding is done with the remaining six electrons. They fi t neatly into a σ bond from two of the p orbitals and two π bonds from the other two pairs.

The electrons in the σ bond lie between the two nuclei, while the electrons in the two π bonds lie in two perpendicular clouds flanking the central σ bond.

Calculating the bond order in N2 is easy—a total of ten bonding electrons and four anti bonding electrons gives a credit of six, or a bond order of three. N2 has a triple-bonded structure. We can’t, however, ignore the electrons that are not involved in bonding: there are eight of them altogether. These non-bonding electrons can be thought of as being localized on each of the N atoms. The four 1s electrons are low-energy inner shell electrons that are not involved in the chemistry of N2, while the four 2s electrons provide the non-bonded lone pairs located one on each N atom. In the structure shown here we have drawn them in: you don’t have to draw lone pairs of every molecule that has them, but sometimes it can be useful to emphasize them—for example if they are taking part in a reaction scheme.

مواضيع ذات صلة

Bond order

Why hydrogen is diatomic but helium is not

Breaking bonds

Molecular orbitals—diatomic molecules

Four short clarifications about orbitals before we go on

The phase of an orbital

s and p orbitals have different shapes

Electrons occupy atomic orbitals

Electrons have quantized energy levels

Atomic emission spectra

Introduction

Looking forward to Chapters 13 and 18