The 118 existing elements (between natural and artificial) keep, to a different extent, a relationship with each other that defines them in families or groups. At first this allowed us to describe the reactivities of the elements, as well as the type of compounds they form; and even better, to predict the properties of those that had not yet been discovered.
Then, as the 20th century passed and physics progressed, these properties were correlated with the electronic structure of the atom. Thus, electrons mark the chemical periodicity with regard to the elements, but not so much for their isotopes and relative stabilities.
Patterns and groups
Chemical periodicity is observed, as expected, in periodic properties. These are characterized by the trend of their values as they are evaluated throughout a period or group of the periodic table.
A zigzag, a mountain range or a steep mountain can be chosen for comparison purposes: with ups and downs. That is, the periodic properties oscillate having minimums and maximums for certain elements. And the relative positions of these elements correspond, brilliantly, with the location in their respective groups.
That is why chemical periodicity is useful to analyze as a function of groups; however, the periods are essential for a complete view of the trend.
You will see with the following examples of periodicity in chemistry, that not only this shines in periodic properties, but also in inorganic and even organic compounds.
Ionization energy, EI, is one of the most outstanding periodic properties. The larger the atom of an element, the easier it will be to remove one of its last electrons; that is to say, those of Valencia. Therefore: atoms with small radii will have large EI, while atoms with large radii will have small EI.
Note, for example, in the image above that the elements Li, Na and K have the lowest EI, which is why they are located in the valleys or bottoms of the graph . Meanwhile, the elements He, Ne and Ar are found at the highest peaks or points, since their EIs are very large in relation to the other elements.
The elements Li, Na and K belong to the group of alkali metals, characterized by their low EI. On the other hand, the elements He, Ne and Ar correspond to the noble gases, with very high EI, because their atoms are the smallest among all the elements for the same period of the periodic table.
Thus, the chemical periodicity indicates that EI decreases going down a group, but increases going one period from left to right.
Hydrides of block p
An example of chemical periodicity outside the periodic properties is seen in the hydrides of the p- block elements .
For example, group 15 is made up of the elements N, P, As, Sb, and Bi. If it is known that ammonia, NH 3 , has nitrogen with an oxidation number of +3, then it can be expected, by simple periodicity, that the remaining elements also have similar hydrides. And indeed it is: NH 3 , PH 3 , AsH 3 , SbH 3 and BiH 3 .
Another similar case occurs with group 16. The hydride of oxygen is H 2 O, water. It is to be expected, again, that the elements S, Se, Te and Po possess hydrides with the same formulas, but with abysmally different properties. And so it is: H 2 S, H 2 Se, H 2 Te and H 2 Po. This is due to chemical periodicity.
Perhaps the most impressive example is the hydrides of carbon, silicon, and germanium. Carbon hydride is CH 4 , and silicon hydride SiH 4 . Germanium, predicted by Mendeleev under the name eka-silicon (Es), must have by periodicity a hydride with the formula GeH 4 ; a prediction that ended up being confirmed after its discovery and subsequent studies.
If fluorine is known to be in the elemental state as a molecule F 2 , then it is to be assumed that the other halogens (Cl, Br, I and At) are also forming diatomic molecules. And so it is, being the molecules Cl 2 , Br 2 and I 2 the best known.
Oxides and sulfides
Similarly, as mentioned with the p- block hydrides , the oxides and sulfides for elements of the same group show a kind of correspondence in their respective chemical formulas. For example, lithium oxide is Li 2 O, the oxides for the other alkali metals or group 1 being: Na 2 O, K 2 O, Rb 2 O and Cs 2 O.
This is due to the fact that all of them have metals with an oxidation number of +1, interacting with an anion O 2- . The same happens with its sulfides: Li 2 S, Na 2 S, etc. In the case of alkaline earth metals or group 2, the formulas of their oxides and sulfides are, respectively: BeO and BeS, MgO and MgS, CaO and CaS, SrO and SrS, BaO and BaS.
This periodicity also applies (in part) to the oxides of the elements of the p block : CO 2 , SiO 2 , GeO 2 , B 2 O 3 , Al 2 O 3 , Ga 2 O 3 , etc. However, for the elements of block d or others of block p , said periodicity becomes more complicated due to the higher possible oxidation numbers for the same element.
For example, copper and silver belong to group 11. One has two oxides: CuO (Cu 2+ ) and Cu 2 O (Cu + ); while the other has just one: AgO (Ag + ).
Hydrocarbons and silanes
Both carbon and silicon have the ability to form CC or Si-Si bonds, respectively. CC bonds are much more stable, so that the structures of hydrocarbons can become disproportionately more numerous and varied than those of their silane counterparts.
This conclusion is due again to chemical periodicity. For example, ethane, CH 3 CH 3 or C 2 H 6 has its homologue disilane, SiH 3 SiH 3 or Si 2 H 6 .