What is the periodic table?

The periodic table of the elements is a tool that allows you to consult the chemical properties of the 118 elements known so far. It is essential when performing stoichiometric calculations, predicting the physical properties of an element, classifying them, and finding periodic properties among all of them.

Atoms become heavier as their nuclei add protons and neutrons, which must also be accompanied by new electrons; otherwise, electroneutrality would not be possible. Thus, some atoms are very light, like hydrogen, and others, super heavy, like oganeson.

To whom is such a heart owed in chemistry? To the scientist Dmitri Mendeleev, who in 1869 (almost 150 years ago) published, after a decade of theoretical studies and experiments, the first periodic table in an attempt to organize the 62 elements known at that time.

For this, Mendeleev relied on chemical properties, while in parallel Lothar Meyer published another periodic table that was organized according to the physical properties of the elements.

The first periodic tables ordered the elements according to their atomic masses. This ordering revealed some periodicity (repetition and similarity) in the chemical properties of the elements; however, the transition elements did not agree with this order, and neither did the noble gases.

For this reason, it was necessary to order the elements considering the atomic number (number of protons), instead of the atomic mass. From here, along with the hard work and contributions of many authors, Mendeleev’s periodic table was refined and completed.

How is the periodic table organized? (Structure and organization)

It can be seen that the pastel of the periodic table has several colors. Each color associates elements with similar chemical properties. There are orange, yellow, blue, purple columns; green squares, and an apple green diagonal.

Note that the cells in the middle columns are grayish in color, so all these elements must have something in common, which is that they are transition metals with half-full d orbitals.

The organization and structure of the periodic table is not arbitrary, but obeys a series of periodic properties and patterns of values ​​determined for the elements. For example, if the metallic character decreases from left to right of the table, a metallic element in the upper right corner cannot be expected.


The elements are arranged in rows or periods depending on the energy level of their orbitals. Before period 4, when the elements succeeded each other in increasing order of atomic mass, it was found that for every eight of them the chemical properties were repeated (John Newlands’ law of octaves).

The transition metals were cast with other non-metallic elements, such as sulfur and phosphorus. For this reason, the entry of quantum physics and electron configurations was vital for the understanding of modern periodic tables.

The orbitals of an energy shell fill up with electrons (and the nuclei of protons and neutrons) as it moves through a period. This energy layer goes hand in hand with the size or atomic radius; therefore, the items in the upper periods are smaller than those below.

H and He are in the first (period) energy level; the first row of grayish squares, in the fourth period; and the row of orange squares, in the sixth period. Note that, although the latter appears to be in the supposed ninth period, it actually belongs to the sixth, just after the yellow box for Ba.


Going through a period it is found that the mass, the number of protons and electrons increase. In the same column or group, although the mass and the protons vary, the number of electrons in the valence shell is the same.

For example, in the first column or group, H has a single electron in the 1s 1 orbital , as does Li (2s 1 ), sodium (3s 1 ), potassium (4s 1 ) and so on until francium (7s 1 ). That number 1 denotes that these elements hardly have a valence electron, and therefore, belong to group 1 (IA). Each item is in different periods.

Not counting green-boxed hydrogen, the elements below it are orange-boxed and are called alkali metals. One more box to the right in any period, is the group or column 2; that is, its elements have two valence electrons.

But when moving one step further to the right, without the knowledge of the d orbitals, one arrives at the boron group (B) or group 13 (IIIA); instead of group 3 (IIIB) or scandium (Sc). Taking into account the filling of the d orbitals, one begins to go through the periods of the grayish squares: the transition metals.

Numbers of protons vs valence electrons

When studying the periodic table a confusion can arise between the atomic number Z or number of total protons in the nucleus, and the number of valence electrons. For example, carbon has a Z = 6, that is, it has six protons and therefore six electrons (otherwise it could not be a neutrally charged atom).

But, of those six electrons, four are of valence . For that reason its electron configuration is [He] 2s 2 2p 2 . [He] denotes the two 1s 2 electrons of the closed shell, and theoretically they do not participate in the formation of chemical bonds .

Also, because carbon has four valence electrons, it “conveniently” is located in group 14 (IVA) of the periodic table.

The elements below carbon (Si, Ge, Sn, Pb and Fl) have higher atomic numbers (and atomic masses); but they all have the four valence electrons in common. This is key to understanding why an item belongs to one group and not another.

Elements of the periodic table

Block s

As just explained, Groups 1 and 2 are characterized by having one or two electrons in s orbitals. These orbitals are of spherical geometry, and as one descends through any of these groups, the elements acquire layers which increase the size of their atoms.

Because they present strong tendencies in their chemical properties and ways of reacting, these elements are organized as the s block. Therefore, the alkali metals and the alkaline earth metals belong to this block. The electronic configuration of the elements of this block is ns (1s, 2s, etc.).

Although the element helium is in the upper right corner of the table, its electronic configuration is 1s 2 and therefore belongs to this block.

Block p

Unlike the s block, the elements of this block have completely filled s orbitals, while their p orbitals continue to be filled with electrons. The electronic configurations of the elements belonging to this block are of the type ns 2 np 1-6 (p orbitals can have one or up to six electrons to fill).

So where on the periodic table is this block located? On the right: the green, purple and blue squares; that is, non-metallic elements and heavy metals, such as bismuth (Bi) and lead (Pb).

Starting with boron, with electronic configuration ns 2 np 1 , the carbon to its right adds another electron: 2s 2 2p 2 . Next, the electron configurations of the other elements of period 2 of block p are: 2s 2 2p 3 (nitrogen), 2s 2 2p 4 (oxygen), 2s 2 2p 5 (fluorine) and 2s 2 2p 6 (neon).

If you go down to the lower periods, you will have the energy level 3: 3s 2 3p 1-6 , and so on until the end of the p block.

Note that the most important thing about this block is that, from period 4, its elements have completely filled d orbitals (blue boxes on the right). In short: block s is on the left of the periodic table, and block p, on the right.

Representative elements

What are the representative elements? They are those that, on the one hand, easily lose electrons, or, on the other, gain them to complete the valence octet. In other words: they are the elements of the s and p blocks.

Their groups were distinguished from the others by a letter A at the end. Thus, there were eight groups: from IA to VIIIA. But currently, the numbering system used in modern periodic tables is Arabic, from 1 to 18, including the transition metals.

For that reason the boron group can be IIIA, or 13 (3 + 10); the carbon group, VAT or 14; and that of noble gases, the last one on the right of the table, VIIIA or 18.

Transition metals

The transition metals are all the elements of the grayish squares. Throughout their periods, their d orbitals are filled, which are five and can therefore have ten electrons. Since they must have ten electrons to fill these orbitals, then there must be ten groups or columns.

Each of these groups in the old numbering system was designated with Roman numerals and a letter B at the end. The first group, that of scandium, was IIIB (3), that of iron, cobalt and nickel VIIIB for having very similar reactivities (8, 9 and 10), and that of zinc IIB (12).

As can be seen, it is much easier to recognize groups by Arabic numbers than by using Roman numerals.

Internal transition metals

As of period 6 of the periodic table, the f orbitals become energetically available. These must be filled first than the d orbitals; and therefore, its elements are usually placed apart so as not to make the table too long.

The last two periods, orange and gray, are the internal transition metals, also called lanthanides (rare earths) and actinides. There are seven f orbitals, which need fourteen electrons to fill, and therefore there must be fourteen groups.

If these groups are added to the periodic table, there will be 32 in total (18 + 14) and there will be a “long” version:

The light pink row corresponds to the lanthanoids, while the dark pink row corresponds to the actinoids. Lanthanum, La with Z = 57, actinium, Ac with Z = 89, and the entire f block belong to the same group as scandium. Why? Because scandium has an nd 1 orbital , which is present in the rest of the lanthanoids and actinoids.

La and Ac have valence configurations 5d 1 6s 2 and 6d 1 7s 2 . As you move to the right through both rows, the 4f and 5f orbitals begin to fill. Once filled, you get to the elements lutetium, Lu, and laurencio, Lr.

Metals and non-metals

Leaving behind the cake of the periodic table, it is more convenient to resort to the one in the upper image, even in its elongated form. At the moment the vast majority of the elements mentioned have been metals.

At room temperature, all metals are solid substances (except mercury, which is liquid) with a silvery-gray color (except for copper and gold). Also, they tend to be hard and shiny; although those of block s are soft and brittle. These elements are characterized by their ease of losing electrons and forming M + cations .

In the case of lanthanoids, they lose the three electrons 5d 1 6s 2 to become trivalent M 3+ cations (such as La 3+ ). Cerium, for its part, is capable of losing four electrons (Ce 4+ ).

On the other hand, non-metallic elements make up the least part of the periodic table. They are gases or solids with covalently linked atoms (like sulfur and phosphorus). All are located in block p; more precisely, in the upper part of it, since descending to the lower periods increases the metallic character (Bi, Pb, Po).

Also, nonmetals instead of losing electrons, you gain them. Thus, they form anions X  with different negative charges: -1 for halogens (group 17), and -2 for chalcogens (group 16, that of oxygen).

Metallic families

Within metals there is an internal classification to differentiate them from each other:

  • Group 1 metals are alkali.
  • Group 2, alkaline earth metals (Mr. Becambara).
  • Group 3 (IIIB) scandium family. This family is made up of scandium, the head of the group, of yttrium Y, lanthanum, actinium, and all lanthanoids and actinoids.
  • Group 4 (IVB), titanium family: Ti, Zr (zirconium), Hf (hafnium) and Rf (rutherfordium). How many valence electrons do they have? The answer is in your group.
  • Group 5 (VB), vanadium family. Group 6 (VIB), chromium family. And so on up to the zinc family, group 12 (IIB).


The metallic character increases from right to left, and from top to bottom. But what is the boundary between these two types of chemical elements? This border is composed of elements known as metalloids, which have characteristics of both metals and non-metals.

Metalloids can be seen on the periodic table in the “ladder” that begins with boron and ends with the radioactive element astatine. These elements are:

  • B: boron.
  • Silicon: Yes.
  • Ge: germanium.
  • As: arsenic.
  • Sb: antimony.
  • Te: tellurium.
  • At: astatine.

Each of these seven elements exhibits intermediate properties, which vary according to chemical environment or temperature. One of these properties is semiconduction, that is, metalloids are semiconductors.


In terrestrial conditions, the gaseous elements are those non-light metals, such as nitrogen, oxygen and fluorine. Also, chlorine, hydrogen and noble gases fall into this classification. Of all of them, the most emblematic are the noble gases, due to their low tendency to react and behave as free atoms.

The latter are found in group 18 of the periodic table and are:

  • Helium, He.
  • Neon, Ne.
  • Argon, Ar.
  • Krypton, Kr.
  • Xenon, Xe.
  • Radon, Rn.
  • And the most recent of all, the synthetic noble gas oganeson, Og.

All noble gases have in common the valence configuration ns 2 np 6 ; that is, they have the entire valence octet.

States of aggregation of elements at other temperatures

The elements are in solid , liquid or gaseous state depending on the temperature and the strength of their interactions. If the temperature of the Earth cooled down to around absolute zero (0K), then all the elements would freeze; except for helium, which would condense.

At this extreme temperature, the rest of the gases would be in the form of ice.

At the other extreme, if the temperature were approximately 6000K, “all” the elements would be in the gaseous state . Under these conditions, you could literally see clouds of gold, silver, lead and other metals.

Uses and applications

The periodic table by itself has always been and will be, a tool for consulting the symbols, atomic masses, structures and other properties of the elements. It is extremely useful when performing stoichiometric calculations, which are the order of the day in many tasks inside and outside the laboratory.

Not only that, but also the periodic table allows you to compare the elements of the same group or period. Thus, one can predict what certain compounds of the elements will be like.

Prediction of oxide formulas

For example, for alkali metal oxides, since they have a single valence electron, and therefore a valence of +1, the formula of their oxides is expected to be of the M 2 O type . This is verified with the oxide of hydrogen, water, H 2 O. Also with the oxides of sodium, Na 2 O, and of potassium, K 2 O.

For the other groups, their oxides must have the general formula M 2 O n , where n is equal to the group number (if the element is from block p, calculate n-10). Thus, carbon, which belongs to group 14, forms CO 2 (C 2 O 4/2 ); sulfur, from group 16, SO 3 (S 2 O 6/2 ); and nitrogen, from group 15, N 2 O 5 .

However, this does not apply to transition metals. This is because iron, although it belongs to group 8, cannot lose 8 electrons but 2 or 3. Therefore, instead of memorizing the formulas, it is more important to pay attention to the valences of each element.

Valences of the elements

The periodic tables (some) show the possible valences for each element. Knowing these, the nomenclature of a compound and its chemical formula can be estimated in advance. Valences, as mentioned earlier, are related to the group number; although it does not apply to all groups.

The valences depend more on the electronic structure of the atoms, and which electrons they can actually lose or gain.

By knowing the number of valence electrons, you can also start with the Lewis structure of a compound from this information. The periodic table therefore enables students and practitioners to sketch structures and make way for a probing of possible molecular geometries and structures.

Digital periodic tables

Today, technology has allowed periodic tables to be more versatile and provide more information available to everyone. Several of them bring striking illustrations of each element, as well as a brief summary of its main uses.

The way you interact with them speeds up your understanding and study. The periodic table should be a tool that is pleasing to the eye, easy to explore, and the most effective method of knowing its chemical elements is to go through it from periods to groups.

Importance of the periodic table

Today, the periodic table is the most important organizing tool in chemistry due to the detailed relationships of its elements. Its use is essential both for students and teachers as well as for researchers and many professionals dedicated to the branch of chemistry and engineering.

Just by looking at the periodic table, you get a huge amount and information quickly and efficiently, such as:

  • Lithium (Li), beryllium (Be), and boron (B) conduct electricity.
  • Lithium is an alkali metal, beryllium is an alkaline earth metal, and boron is a nonmetal.
  • Lithium is the best conductor of the three named, followed by beryllium and, lastly, boron (semiconductor).

Thus, by locating these elements in the periodic table, their tendency to electrical conductivity can be instantly concluded.

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