Fluorine has five electrons in its p-orbital. This is not a sphere-shaped harmonic, and so five p-electrons cannot achieve a stable electron symmetry around a spherical core. Fluorine therefore cannot simply add its five p-electrons on top of the same (2s2) configuration that beryllium has, as shown here, because it would not be a stable configuration:
The asymmetry therefore causes fluorine to hybridize its 2s and 2p electrons in order to achieve tetrahedral symmetry. Its sp3 hybrid orbitals feature three di-electrons and one unpaired electron, rendering it extremely (and dangerously) reactive in search of that final electron-pairing. One more electron will give it a full 2nd shell, like neon, and that is a very attractive state for the atom. In addition, a high effective nuclear charge gives fluorine the highest electronegativity in its row, and because it is the smallest of the Group VII elements, its electronegativity is also the highest on the periodic table. (The wireframe indicates the boundary of the n=2 shell.)
The outer spheres above simply indicate the directions of maximum electron density. The orbitals themselves will be more like spherical tetrahedra that can only occupy volume within their shell. The entire shell will be filled with electron density. It will be highest at the center of the face of each orbital and will decrease toward the nodal regions that divide the orbitals, where electron density will be lowest. (This includes the inner and outer faces, as well as the sides of each orbital.) In the case of fluorine, the three orbitals containing di-electrons will each occupy slightly more volume than the one containing the unpaired electron.
Fluorine is so eager to obtain an extra electron to fill its second shell that it can bond with just about any atom on the periodic table, even several of the usually-unreactive noble gases, forcing them to donate electrons into that bond. Fluorine can therefore make a single covalent bond or it can gain an electron in ionic interaction in order to reach the stability of a full 2nd shell. This is the same electron configuration as the 2s22p6 noble gas configuration of neon — a multi-di-electron state with two concentric full shells. That is why fluorine forms a 1– ionic state. The negative ion is larger than the neutral atom because electrons now outnumber protons by one. This results in a lower effective nuclear charge — a lower average attraction by the nucleus on each electron.
The atoms of the second row of the periodic table are small, making the nuclei relatively much closer to their surface than in larger atoms. Fluorine not only lacks one electron to complete its stable tetrahedral dielectron resonance, but the high effective nuclear charge of its (shallow) nucleus adds even more strength to fluorine’s strong attraction on nearby electrons. This makes it so highly reactive that it will react with (in order to pull an electron from) every other element on the periodic table except the two smallest noble gas atoms, helium and neon.
If fluorine gas is simply passed over carbon, for example, the carbon will spontaneously combust in it. Although hydrofluoric acid (HF) is a weak acid, it cannot be stored in a glass container because it will degrade the glass.
RETURN to the Periodic Table