PERIODIC PROPERTIES OF ELEMENTS AND THEIR CHARACTERISTICS - CHEMISTRY - 2026

The periodic properties of the elements are those that define their physical and chemical behavior from an atomic perspective, and whose magnitudes, in addition to the atomic number, allow a classification of atoms.

Of all the properties, these are characterized, as their name indicates, for being periodic; that is, if the periodic table is studied, it will be possible to certify that its magnitudes obey a trend that coincides and is repeated with the ordering of the elements in periods (rows) and groups (columns).

Intrinsic periodicity of a part of the elements of the periodic table. Source: Gabriel Bolívar.

For example, if a period is traversed and a periodic property decreases in magnitude with each element, the same will happen in all periods. On the other hand, if going down one group or column increases its magnitude, the same can be expected to happen for the other groups.

And so, its variations are repeated and show a simple tendency that agrees with the ordering of the elements by their atomic numbers. These properties are directly responsible for the metallic or non-metallic character of the elements, as well as their reactivities, which has helped to classify them in greater depth.

If for a moment the identity of the elements were unknown and they were seen as strange "spheres", the periodic table could be rebuilt (with a lot of work) using these properties.

In this way, the supposed spheres would acquire colors that would allow them to be differentiated from one another in groups (upper image). Knowing their electronic characteristics, they could be organized in periods, and the groups would reveal those that have the same number of valence electrons.

Learning and reasoning about periodic properties is the same as knowing why elements react in one way or another; is to know why the metallic elements are in certain regions of the table, and the non-metallic elements in another.

What are the periodic properties and their characteristics

-Atomic radio

When observing the spheres in the image, the first thing that can be noticed is that they are not all the same size. Some are more voluminous than others. If you look more closely, you will find that these sizes vary according to a pattern: in one period it decreases from left to right, and in a group it increases from top to bottom.

The above can also be said this way: the atomic radius decreases towards the groups or columns on the right, and increases in the lower periods or rows. This being so, the atomic radius is the first periodic property, since its variations follow a pattern within the elements.

Nuclear charge vs electrons

What is the cause of this pattern? In a period the electrons of the atom occupy the same energy level, which is related to the distance that separates them from the nucleus. When we move from one group to another (which is the same as going through the period to the right), the nucleus adds both electrons and protons within the same energy level.

Therefore, the electrons cannot occupy further distances from the nucleus, which increases its positive charge because it has more protons. Consequently, the electrons experience a greater force of attraction towards the nucleus, attracting them more and more as the number of protons increases.

That is why the elements on the far right of the periodic table (yellow and turquoise columns) have the smallest atomic radii.

On the other hand, when you "jump" from one period to another (which is the same as saying that you descend through a group), the new energy levels enabled allow the electrons to occupy more distant spaces from the nucleus. Being further away, the nucleus (with more protons) attracts them with less force; and the atomic radii therefore increase.

Ionic radii

Ionic radii follow a similar pattern to atomic radii; However, these do not depend so much on the nucleus, but on how many or fewer electrons the atom has with respect to its neutral state.

The cations (Na ⁺, Ca ²⁺, Al ³⁺, Be ²⁺, Fe ³⁺) exhibit a positive charge because they have lost one or more electrons, and therefore, the nucleus attracts them with greater force since there are fewer repulsions. between them. The result: cations are smaller than the atoms from which they are derived.

And for anions (O ^2-, F ^-, S ^2-, I ^-) on the contrary, they exhibit negative charge because they have one or more electrons in excess, increasing their repulsions to each other above the attraction exerted by the nucleus. The result: the anions are larger than the atoms from which they are derived (image below).

Variation of ionic radii with respect to the neutral atom. Source: Gabriel Bolívar.

It is appreciated that the 2- anion is the largest of all, and the 2+ cation the smallest. The radii increase when the atom is negatively charged, and contracts when it is positively charged.

-Electronegativity

When elements have small atomic radii, not only are their electrons attracted very strongly, but also electrons from neighboring atoms when they form a chemical bond. This tendency to attract electrons from other atoms within a compound is known as electronegativity.

Just because an atom is small does not mean that it will be more electronegative. If so, the elements helium and hydrogen would be the most electronegative atoms. Helium, as far as science has shown, does not form a covalent bond of any kind; and hydrogen only has a single proton in the nucleus.

When the atomic radii are large, the nuclei are not strong enough to attract electrons from other atoms; therefore, the most electronegative elements are those with a small atomic radius and a greater number of protons.

Again, those that fulfill these characteristics perfectly are the non-metallic elements of the p block of the periodic table; These are those belonging to group 16 or oxygen (O, S, Se, Te, Po), and group 17 or fluorine (F, Cl, Br, I, At).

Trend

According to all that has been said, the most electronegative elements are located especially in the upper right corner of the periodic table; having fluorine as the element that heads the list of the most electronegative.

Why? Without resorting to electronegativity scales (Pauling, Mulliken, etc.), although fluorine is larger than neon (the noble gas of its period), the former can form bonds while the latter cannot. Also, for its small size, its nucleus has many protons, and where the fluorine is, there will be a dipole moment.

-Metallic character

If an element has an atomic radius compared to those of its same period, and is also not very electronegative, then it is a metal, and has a high metallic character.

If we return to the main image, the reddish and greenish spheres, like the grayish ones, correspond to metallic elements. Metals have unique characteristics, and from here the periodic properties begin to intertwine with the physical and macroscopic properties of matter.

The elements with high metallic character are characterized by their relatively large atoms, easy to lose electrons since the nuclei can hardly attract them towards them.

As a result, they are easily oxidized or lost electrons to form cations, M ⁺; this does not mean that all cations are metallic.

Trend

At this point you can predict how the metallic character varies in the periodic table. If it is known that metals have large metallic radii, and that they are also few electronegative, it must be expected that the heaviest elements (the lower periods) are the most metallic; and the lightest elements (the upper periods), the least metallic.

Also, the metallic character decreases the more electronegative the element becomes. This means that going through the periods and groups to the right of the periodic table, in their upper periods, they will find the less metallic elements.

Therefore, the metallic character increases descending through a group, and decreases from left to right in the same period. Among the metallic elements we have: Na (sodium), Li (lithium), Mg (magnesium), Ba (barium), Ag (silver), Au (gold), Po (polonium), Pb (lead), Cd (cadmium), Al (aluminum), etc.

-Ionization energy

If an atom has a large atomic radius, it is to be expected that its nucleus will not hold electrons in the outermost shells trapped with considerable force. Consequently, removing them from the atom in the gas phase (individualized) will not require much energy; that is, the ionization energy, EI, necessary to remove an electron from them.

EI is also equivalent to saying that it is the energy that must be supplied to overcome the attractive force of the nucleus of an atom or gaseous ion on its outermost electron. The smaller the atom and the more electronegative, the lower its EI; this is your trend.

The following equation illustrates an example:

Na (g) => Na ⁺ (g) + e ^-

The EI needed to achieve this is not that great compared to the second ionization:

Na ⁺ (g) => Na ²⁺ (g) + e ^-

Since in Na ⁺ positive charges predominate and the ion is smaller than the neutral atom. Consequently, the nucleus of Na ⁺ attracts electrons with a much greater force, requiring a much larger EI.

-Electronic affinity

And finally, there is the periodic property of electronic affinity. This is the energetic tendency of the atom of an element in the gas phase to accept an electron. If the atom is small and has a nucleus with a large attractive force, it will be easy for it to accept the electron, forming a stable anion.

The more stable the anion is with respect to its neutral atom, the greater its electron affinity. However, repulsions between the electrons themselves also come into play.

Nitrogen, for example, has a higher electron affinity than oxygen. This is because its three 2p electrons are unpaired and repel each other and the incoming electron less; while in oxygen, there is a pair of paired electrons that exert greater electronic repulsion; and in fluorine, there are two pairs.

It is for this reason that the trend in electronic affinities is said to normalize from the third period of the periodic table.

References

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