PERIODIC TABLE OF ELEMENTS: HISTORY, STRUCTURE, ELEMENTS - CHEMISTRY

The periodic table of the elements is a tool that allows 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.

To do 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.

Initially, the table contained “empty spaces”, the elements of which had not been known in those years. However, Mendeleyev was able to predict several of its properties with appreciable accuracy. Some of these elements were: germanium (which he called eka-silicon) and gallium (eka-aluminum).

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.

History of the periodic table

Elements

The use of elements as the basis to describe the environment (more precisely, nature) has been used since ancient times. However, at that time they were referred to as the phases and states of matter, and not in the way in which they are referred to from the Middle Ages.

The ancient Greeks had the belief that the planet we inhabit was made up of the four fundamental elements: fire, earth, water and air.

On the other hand, in ancient China the number of elements was five and, unlike the Greeks, these excluded air and included metal and wood.

The first scientific discovery was made in 1669 by the German Henning Brand, who discovered phosphorus; as of that date, all subsequent items were recorded.

It is worth clarifying that some elements such as gold and copper were already known before phosphorus; the difference is that they were never registered.

Symbology

The alchemists (forerunners of today's chemists) gave names to the elements in relation to the constellations, their discoverers and the places where they were discovered.

In 1808 Dalton proposed a series of drawings (symbols) to represent the elements. Later, this notation system was replaced by that of Jhon Berzelius (used to date), since Dalton's model became more complicated as new elements appeared.

Scheme evolution

The first attempts to create a map that organized the information of the chemical elements occurred in the 19th century with the Döbereiner Triads (1817).

Over the years, new elements were found, giving rise to new organizational models until reaching the one currently used.

Telluric screw of Chancourtois (1862)

Alexandré-Émile Béguyer de Chancourtois designed a paper helix showing a graph of spirals (telluric screw).

In this system the elements are ordered in increasing order with respect to their atomic weights. Similar items are vertically aligned.

Octaves of Newlands (1865)

Continuing with Döbereiner's work, the British John Alexander Reina Newlands arranged the chemical elements in increasing order with respect to atomic weights, noting that every seven elements had similarities in their properties (hydrogen is not included).

Mendeleev's Table (1869)

Mendeleev arranged the chemical elements in increasing order with respect to atomic weight, placing in the same column those whose properties were similar. He left gaps in his model of the periodic table anticipating the appearance of new elements in the future (in addition to predicting the properties it should have).

The noble gases do not appear in Mendeleev's table, since they had not yet been discovered. Furthermore, Mendeleiv did not consider hydrogen.

Moseley's periodic table (current periodic table) - 1913

Henry Gwyn Jeffreys Moseley proposed to order the chemical elements of the periodic table according to their atomic number; that is, based on their number of protons.

Moseley enunciated the "Periodic Law" in 1913: "When elements are arranged in order of their atomic numbers, their physical and chemical properties show periodic trends."

Thus, each horizontal row or period shows one type of relationship, and each column or group shows another.

How is it organized? (Structure and organization)

It can be seen that the periodic table pastel 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.

In the same way, the elements of the purple squares, although they go from gaseous substances, from a reddish liquid to solid black-purple (iodine) and silver-gray (astatine), it is their chemical properties that make them congeners. These properties are governed by the electronic structures of its atoms.

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.

Periods

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 repeated themselves (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 travels over a period of time. 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.

Groups

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 consideration the filling of the d orbitals, one begins to go through the periods of the grayish squares: the transition metals.

Proton numbers vs valence electrons

When studying the periodic table, a confusion may 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 2s ² 2p ². 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, "conveniently" it 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 block p.

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:

Source: By Sandbh, from Wikimedia Commons

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 are usually hard and shiny; although those of block s are soft and fragile. 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 (such as 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:

-The metals of group 1 are alkaline

-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).

Metalloids

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.

Gases

In terrestrial conditions, the gaseous elements are those light non-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:

-Helio, 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, even though 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 above, are related to the group number; although it does not apply to all groups.

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

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 allows students and professionals to sketch structures and to make way for a probing of possible geometries and molecular 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 their 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 non-metal.

- 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.

References

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PERIODIC TABLE OF ELEMENTS: HISTORY, STRUCTURE, ELEMENTS - CHEMISTRY - 2025