So far, so good. Now, lithium. The lowest energy 1s orbital around the Li nucleus can contain two electrons, but two only, so the third electron has to go into a higher energy orbital—one of the energy levels whose existence was inferred from atomic absorption spectroscopy. You can think of this orbital as the three-dimensional equivalent of the first harmonic of the gui tar string. Like the vibration of the string, this next orbital has a node. On the string the node was the point where no motion was observed. In an atom, a node is a point where the electron can never be found—a void separating the two parts of the orbital. For the orbital containing the third electron of the Li atom, this node is spherical—it divides the orbital into two parts which nestle one within another like the layers of an onion or the stone inside a peach. We call this orbital the 2s orbital—’2’ because we have moved up to an orbital with a node (like the first harmonic) and ‘s’ because the orbital is still spherical. The ‘s’ did not originally stand for ‘spherical’, but as all ‘s’ orbitals are spherical it’s fine to remember it that way.

In a lithium atom the 1s orbital, close to the nucleus, is occupied by two electrons, while the 2s orbital, further from the nucleus, contains one. In beryllium, there is a second electron in the 2s orbital. As before the energy levels will change as the nuclear charge increases, so the orbital occupancy in Li and Be can be represented as shown below.

When we get to boron, something a little different happens. It turns out that for an orbital with one node (such as the 2s orbital), the node does not have to be spherical. The node can alternatively be a plane. This alternative arrangement for an orbital with a single planar node gives us a new type of orbital, the 2p orbital. A 2p orbital looks something like the picture on the left below, in ‘smeared out’ form. It is often represented as the propeller shape in the mid-dle, and it is conventionally drawn as the shape shown in the diagram to the right.

Unlike the 1s or 2s orbitals, the 2p orbital is directional—it points along an axis, and in three dimensions there are three possible orientations for the axis, each of which gives rise to a new 2p orbital (which we can call 2px, 2py and 2pz if we need to).

orbital, with its spherical node. Boron atoms therefore have two electrons in the 1s orbital, two in the 2s orbital, and just one in one of the 2p orbitals. The orbital occupancy is shown in the energy level diagram on the left. You can imagine what shape each of the orbitals has: we won’t need to show a picture of them all superimposed. The next element, carbon, with one more (a sixth) electron, seems to have a choice—it can put its sixth electron paired with the fifth one, in the same 2p orbital, or it can put it into a new 2p orbital, with both electrons unpaired. In fact, it chooses the latter: electrons are negatively charged and repel one another, so if there is a choice of equal energy orbitals, they occupy different orbitals singly until they are forced to start pairing up. The repulsion is never enough to force an electron to occupy a higher energy orbital, but when the orbitals are other wise of identical energy, this is what happens. Not surprisingly therefore, the orbitals of atoms of the remainder of the elements of the fi rst row of the periodic table are occupied as shown below. All the while the entire set of orbitals is going down in energy because the nucleus is attracting the electrons more strongly, but otherwise it is a simple matter of filling up the 2p orbitals fi rst singly and then doubly. With the ten electrons of neon, all the orbitals with one node are filled, and we say that neon has a ‘closed shell’. A ‘shell’ is a group of orbitals of similar energy all with the same number of nodes (in this case all called ‘2’ something—2s or 2p).