You have learned that atoms, in general, are not very stable. In fact, there are only six known chemically stable atoms: the noble gases. The noble gas atoms are chemically stable because of their electron configurations. Each of them has a “full” set of electrons in their outermost shell.
But wait a minute. The helium and neon atoms both have full outer electron shells… the helium atom has two electrons in its first (outer) shell and the neon atom has eight electrons in its second (outer) shell… but the argon atom only has eight electrons in its third (outer) shell when the third shell can hold up to eighteen electrons. How is the argon atom’s outer shell anywhere close to full?
The reason why is because the Bohr model of an atom, based on electron shells, is not accurate. Electrons do not occupy electron shells; they occupy atomic orbitals. The next element after argon is potassium (K). The potassium atom has one more proton and one more electron than the argon atom. The Bohr model predicts that this nineteenth electron will go into the third electron shell. But in reality, the nineteenth electron goes into the s-orbital in the fourth electron shell, even though the third shell is not completely filled yet. The remaining orbitals in the third shell only get filled once the s-orbital in the fourth shell has been filled first. (Bohr diagrams are only accurate up to Argon.)
Although the third electron shell can hold 18 electrons, the fourth electron shell can hold 32 electrons, the fifth electron shell can hold 50 electrons, and the sixth electron shell can hold 72 electrons, once each of those electron shells gets filled with 8 electrons, the next two electrons will go into the next higher electron shell. This means that the outermost shell of an atom can only have between 1 and 8 electrons in it before a new outer shell gets started, and this shell becomes an inner shell instead. The electrons in the outermost shell of an atom are called valence electrons.
If an atom only has one electron, then all of the atomic orbitals of an electron shell are equivalent and have the same energy, and it does not matter which atomic orbital an electron fills. But when an atom has more than one electron, the geometry of the atomic orbitals and the positions of other electrons do affect the energy of an atomic orbital and the order in which they are filled.
The first electron shell has one atomic orbital: an s-orbital. All s-orbitals are shaped like spheres. The second electron shell has four atomic orbitals: one s-orbital and three p-orbitals. The p-orbitals are shaped like dumbbells in order to maximize the distance between the electrons. Maximizing the distance between electrons makes an atom more stable because electrons are negatively charged and repel each other.
The two electrons in the px-orbital spend most of their time in regions to the left and right of the atomic nucleus; the two electrons in the py-orbital spend most of their time in regions above and below the atomic nucleus; and the two electrons in the pz-orbital spend most of their time in regions in front of and behind the atomic nucleus. The geometry of these orbitals spread the electrons out as much as possible, but there is still some overlap, which lowers the stability and increases the energy of the p-orbitals slightly. This is why, in reality, the s-orbital is filled before the p-orbitals.
The third electron shell has nine atomic orbitals: one s-orbital, three p-orbitals, and five d-orbitals. The d-orbitals have even more complicated and stretched out shapes in order to maximize the distance between the electrons, but there is so much overlap that the s-orbital in the fourth electron shell actually gets filled before the five d-orbitals in the third electron shell start getting filled.
This table shows the order in which the atomic orbitals get filled. 1s is the s-orbital in the first electron shell. 2s is the s-orbital in the second electron shell. 2p is the p-orbitals in the second electron shell.
Because most atoms are not very stable, they tend to chemically bond with other atoms to form stable molecules. The way that many atoms bond can be predicted by the octet rule. The octet rule states that atoms tend to lose, accept, or share electrons until they have an electron configuration with a “full” outer shell (full, in this case, meaning two electrons if the outermost shell is the first shell, and eight electrons if the outermost shell is a second or higher shell). This is the same electron configuration that noble gas atoms have, which are the only chemically stable atoms in the periodic table. The octet rule does not always work since there are so many variables that can affect an atom’s stability. In general, the octet rule does not apply to transition or inner transition metals.
As you have seen, Bohr diagrams can be a little misleading. They are also not terribly useful for predicting the chemical properties of atoms. Chemical properties describe how an atom or molecule reacts in the presence of other atoms or molecules. Atoms react by losing, accepting, or sharing electrons, and this almost always involves an atom’s outermost (valence) electrons, not the electrons in the inner shells. A Lewis structure is a diagram that highlights an atom’s valence electrons. It can be very useful for predicting how atoms will react to each other.
Some people fill in the dots (valence electrons) of a Lewis structure in different orders. There does not seem to be one right way to do it, but this is the way that I prefer. Start by filling in the s-orbital electrons to the right of the atomic symbol, and then fill in the p-orbital electrons one-by-one, going in the counterclockwise direction around the symbol.
If you were to draw Lewis structures for all of the elements in the periodic table, you would find an interesting pattern. All of the elements in the seventeenth column of the periodic table have the same Lewis structures.
These elements are collectively known as the halogens or the fluorine group. Because they have seven valence electrons in the same electron configuration, they have similar chemical properties. The halogens are highly reactive, and will bond with other atoms by accepting or sharing one electron to fill their valence shells.
At the other end of the periodic table are the alkali metals or the lithium group. Because they have one valence electron in the same electron configuration, the alkali metals also have similar chemical properties. They are highly reactive, and will bond with other atoms by losing their one valence electron. This makes the next lower (and full) electron shell the outermost, valence shell. While hydrogen is technically part of this group, it has some different chemical properties. This is because, in addition to losing its electron, a hydrogen atom can also accept or share an electron to fill its valence shell.