Knowledge database: 1.10 Electrons

To understand why and how the atoms of different elements connect, it is necessary to know more about the atomic structure. If you look at the periodic table of elements (PTE), you can see that each atom has its (atomic) number. This number tells how many protons an atom has in its nucleus, and how many electrons there are in the electrone layer. Hydrogen has one proton and one electron, which is also located in his last (and only) orbital - an energy level containing electrons. The next atom in line, helium, has 2 electrons, which are still in a single energy level (orbital). Orbitals can, for simplicity, be imagined as rings around the nucles in which the electrons circulate. In these rings 2 electrons can be placed (the s orbital). Therefore, with the next element in the row, lithium, the first ring is filled, and the second ring (also an s orbital), contains only 1 electron. To avoid describing this element by element, the scheme by which electrons are placed in orbitals can be seen below.

Electron placement scheme

As we see, there are s, p, d and f orbitals. The s orbitals hold 2 electrons, p orbitals hold 6, d orbitals 10, and f orbital holds 14 electrons. When you fill the 1s orbital, you go on to the 2s orbital. After the 2s orbital is full, next comes the 2p orbitals (which unlike the s orbitals, hold 6 electrons). It is also necessary to describe exactly how the electrons are filled into the orbitals (picture below).

Electron positioning method

Each square has "room" for 2 electrons. Since the s orbital receives only 2 electrons, only one box is shown. P orbitals are represented with 3 squares because they can recieve 6 electrons, etc. It is also very important to mention that the electrons are first filled by introducing 1 electron per square, and when all squares have one electron, only then electrons are introduced to squares where 1 electron is already located. As an example, take a look at the d orbitals in the picture above. We have 6 electrons which have filled d orbitals squares. The first 5 are deployed so that one is placed in each box, and then we place the sixth electron in the first box along with the electron which we have previously placed there. By following these rules you can fill the orbitals of any atom/element. This is important because by knowing the number of free, unmatched electrons (those that stand alone in a square of some of the last orbitals) you can determine the oxidation numbers of an atom. More on this in the next section.