Periodic Table

The **periodic table**, also known as the **periodic table of** (**the**) (**chemical**) **elements**, is a tabular display of the chemical elements. It is widely used in chemistry, physics, and other sciences, and is generally seen as an icon of chemistry. It is a graphic formulation of the periodic law, which states that the properties of the chemical elements exhibit a periodic dependence on their atomic numbers.

The table is divided into four roughly rectangular areas called blocks. The rows of the table are called periods, and the columns are called groups. Elements from the same column group of the periodic table show similar chemical characteristics. Trends run through the periodic table, with non-metallic character (keeping their own electrons) increasing from left to right across a period, and from down to up across a group, and metallic character (surrendering electrons to other atoms) increasing in the opposite direction. The underlying reason for these trends is electron configurations of atoms.

The first periodic table to become generally accepted was that of the Russian chemist Dmitri Mendeleev in 1869: he formulated the periodic law as a dependence of chemical properties on atomic mass. Because not all elements were then known, there were gaps in his periodic table, and Mendeleev successfully used the periodic law to predict properties of some of the missing elements. The periodic law was recognized as a fundamental discovery in the late 19th century, and it was explained with the discovery of the atomic number and pioneering work in quantum mechanics of the early 20th century that illuminated the internal structure of the atom. With Glenn T. Seaborg's 1945 discovery that the actinides were in fact f-block rather than d-block elements, a recognisably modern form of the table was reached. The periodic table and law are now a central and indispensable part of modern chemistry.

The periodic table continues to evolve with the progress of science. In nature, only elements up to atomic number 94 exist; to go further, it was necessary to synthesise new elements in the laboratory. Today, all the first 118 elements are known, completing the first seven rows of the table, but chemical characterisation is still needed for the heaviest elements to confirm that their properties match their positions. It is not yet known how far the table will stretch beyond these seven rows and whether the patterns of the known part of the table will continue into this unknown region. Some scientific discussion also continues regarding whether some elements are correctly positioned in today's table. Many alternative representations of the periodic law exist, and there is some discussion as to whether or not there is an optimal form of the periodic table.

## Overview

The smallest constituents of all normal matter are known as atoms. Atoms are extremely small, being about one ten-billionth of a meter across; thus their internal structure is governed by quantum mechanics. Atoms consist of a small positively charged nucleus, made of positively charged protons and uncharged neutrons, surrounded by a cloud of negatively charged electrons; the charges cancel out, so atoms are neutral. Electrons participate in chemical reactions, but the nucleus does not. When atoms participate in chemical reactions, they may gain or lose electrons to form positively- or negatively-charged ions; or they may share electrons with each other instead.

Atoms can be subdivided into different types based on the number of protons (and thus also electrons) they have. This is called the atomic number, often symbolised *Z* as German for number is *Zahl*. Each distinct atomic number therefore corresponds to a class of atom: these classes are called the chemical elements. The chemical elements are what the periodic table classifies and organises. Hydrogen is the element with atomic number 1; helium, atomic number 2; lithium, atomic number 3; and so on. Each of these names can be further abbreviated by a one- or two-letter chemical symbol; those for hydrogen, helium, and lithium are respectively H, He, and Li. Neutrons do not affect the atom's chemical identity, but do affect its weight. Atoms with the same number of protons but different numbers of neutrons are called isotopes of the same chemical element. Naturally occurring elements usually occur as mixes of different isotopes; since each isotope usually occurs with a characteristic abundance, naturally occurring elements have well-defined atomic weights, defined as the average mass of a naturally occurring atom of that element.

Today, 118 elements are known, the first 94 of which occur in nature. Of the 94 natural elements, eighty are stable; three more (bismuth, thorium, and uranium) undergo radioactive decay, but so slowly that large quantities survive from the formation of the Earth; and eleven more decay quickly enough that their continued trace occurrence rests on being constantly regenerated as intermediate products of the decay of thorium and uranium. The 83 elements that have survived from the formation of the Earth are called primordial. All 24 known artificial elements are radioactive.

The periodic table is a graphic description of the periodic law, which states that the properties and atomic structures of the chemical elements are a periodic function of their atomic number. Elements are placed in the periodic table by their electron configurations, which exhibit periodic recurrences that explain the trends of properties across the periodic table.

An electron can be thought of as inhabiting an atomic orbital, which characterises the probability it can be found in any particular region of the atom. Their energies are quantised, which is to say that they can only take discrete values. Furthermore, electrons obey the Pauli exclusion principle: different electrons must always be in different states. This allows classification of the possible states an electron can take in various energy levels known as shells, divided into individual subshells, which each contain a certain kind of orbital. Each orbital can contain up to two electrons: they are distinguished by a quantity known as spin, which can be up or down. Electrons arrange themselves in the atom in such a way that the total energy they have is minimised, so they occupy the lowest-energy orbitals available unless energy has been supplied. Only the outermost electrons (so-called valence electrons) have enough energy to break free of the nucleus and participate in chemical reactions with other atoms. The others are called core electrons.

Elements are known with up to the first seven shells occupied. The first shell contains only one orbital, a spherical s orbital. As it is in the first shell, this is called the 1s orbital. This can hold up to two electrons. The second shell similarly contains a 2s orbital, but it also contains three dumbbell-shaped p orbitals, and can thus fill up to eight electrons (2×1 + 2×3 = 8). The third shell contains one 3s orbital, three 3p orbitals, and five 3d orbitals, and thus has a capacity of 2×1 + 2×3 + 2×5 = 18. The fourth shell contains one 4s orbital, three 4p orbitals, five 4d orbitals, and seven 4f orbitals, thus leading to a capacity of 2×1 + 2×3 + 2×5 + 2×7 = 32. Higher shells contain more types of orbitals that continue the pattern, but such types of orbitals are not filled in the known elements. The subshell types are characterised by the quantum numbers. Four numbers describe an electron in an atom completely: the principal quantum number *n* (the shell), the azimuthal quantum number ℓ (the orbital type), the magnetic quantum number *m*ℓ (which of the orbitals of a certain type it is in), and the spin quantum number *s*.

The sequence in which the orbitals are filled is given by the Aufbau principle, also known as the Madelung or Klechkovsky rule. The shells overlap in energies, creating a sequence that goes approximately as follows:

1s ≪ 2s < 2p ≪ 3s < 3p ≪ 4s < 3d < 4p ≪ 5s < 4d < 5p ≪ 6s < 4f < 5d < 6p ≪ 7s < 5f < 6d < 7p ≪ ...
Here the sign ≪ means "much less than" as opposed to < meaning just "less than". Phrased differently, electrons enter orbitals in order of increasing *n* + ℓ, and if two orbitals are available with the same value of *n* + ℓ, the one with lower *n* is occupied first.

The overlaps get quite close at the point where the d-orbitals enter the picture, and the order can shift slightly with atomic number and atomic charge.

Starting from the simplest atom, this lets us build up the periodic table one at a time in order of atomic number, by considering the cases of single atoms. In hydrogen, there is only one electron, which must go in the lowest-energy orbital 1s. This configuration is thus written 1s1. Helium adds a second electron, which also goes into 1s and fills the first shell completely.

The third element, lithium, has no more space in the first shell. Its third electron must thus enter the 2s subshell, giving a 1s2 2s1 configuration. The 2s electron is lithium's only valence electron, as the 1s orbital is now too close to the nucleus to participate chemically. The 2s subshell is completed by the next element beryllium (1s2 2s2). The following elements then proceed to fill up the p-orbitals. Boron (1s2 2s2 2p1) puts its new electron in a 2p orbital; carbon (1s2 2s2 2p2) fills a second 2p orbital; and with nitrogen (1s2 2s2 2p3) all three 2p orbitals become singly occupied. This is consistent with Hund's rule, which states that atoms will prefer to singly occupy each orbital of the same type before filling them with the second electron. Oxygen (1s2 2s2 2p4), fluorine (1s2 2s2 2p5), and neon (1s2 2s2 2p6) then complete the already singly filled 2p orbitals; the last of these fills the second shell completely.

Starting from element 11, sodium, there is no more space in the second shell, which from here on is a core shell just like the first. Thus the eleventh electron enters the 3s orbital instead. The 1s2 2s2 2p6 core is abbreviated [Ne], as it is identical to the electron configuration of neon: thus sodium's configuration is written [Ne] 3s1. Magnesium ([Ne] 3s2) finishes this 3s orbital, and from then on the six elements aluminium, silicon, phosphorus, sulfur, chlorine, and argon fill the three 3p orbitals ([Ne] 3s2 3p1 through [Ne] 3s2 3p6). This creates an analogous series in which the outer shell structures of sodium through argon are exactly analogous to those of lithium through neon, and is the basis for chemical periodicity that the periodic table illustrates: at regular but changing intervals of atomic numbers, the properties of the chemical elements approximately repeat.

The first eighteen elements can thus be arranged as the start of a periodic table. Elements in the same column have the same number of outer electrons and analogous outer electron configurations: these columns are called groups. The single exception is helium, which has two outer electrons like beryllium and magnesium, but is placed with neon and argon to emphasise that its outer shell is full. There are eight columns in this periodic table fragment, corresponding to at most eight outer electrons. A row begins when a new shell begins filling; these rows are called the periods. Finally, the colouring illustrates the blocks: the elements in the s-block (coloured red) are filling s-orbitals, while those in the p-block (coloured yellow) are filling p-orbitals.

Starting the next row, for potassium and calcium the 4s orbital is the lowest in energy, and therefore they fill it. Potassium adds one electron to the 4s shell ([Ar] 4s1), and calcium then completes it ([Ar] 4s2). However, starting from scandium the 3d orbital becomes the next highest in energy. The 4s and 3d orbitals are of approximately the same energy and they compete for filling the electrons, and so the occupation is not quite consistently filling the 3d orbitals one at a time. The precise energy ordering of 3d and 4s changes along the row, and also changes depending on how many electrons are removed from the atom. For example, due to the repulsion between the 3d electrons and the 4s ones, at chromium the 4s energy level becomes slightly higher than 3d, and so it becomes more profitable to have a [Ar] 3d5 4s1 configuration than an [Ar] 3d4 4s2 one. A similar anomaly occurs at copper. These are violations of the Madelung rule. Such anomalies however do not have any chemical significance, as the various configurations are so close in energy to each other that the presence of a nearby atom can shift the balance. The periodic table therefore ignores these and considers only idealised configurations.

At zinc, the 3d orbitals are completely filled with a total of ten electrons. Next come the 4p orbitals completing the row, which are filled progressively by gallium through krypton, in a manner totally analogous to the previous p-block elements. From gallium onwards, the 3d orbitals form part of the electronic core, and no longer participate in chemistry. The s- and p-block elements, which fill their outer shells, are called main-group elements; the d-block elements (coloured blue below), which fill an inner shell, are called transition elements (or transition metals, since they are all metals).

As 5s fills before 4d, which fills before 5p, the fifth row has exactly the same structure as the fourth (though again there are some anomalies when filling the 4d orbitals).

The sixth row of the table likewise starts with two s-block elements: caesium and barium. After this, the first f-block elements (coloured green below) begin to appear, starting with lanthanum. These are sometimes termed inner transition elements. As there are now not only 4f but also 5d and 6s subshells at similar energies, competition occurs once again with many irregular configurations; this has resulted in some dispute about where exactly the f-block is supposed to begin, but most who study the matter agree that it starts at lanthanum in accordance with the Aufbau principle. Even though lanthanum does not itself fill the 4f orbital due to repulsion between electrons, its 4f orbitals are low enough in energy to participate in chemistry. At ytterbium, the seven 4f orbitals are completely filled with fourteen electrons; thereafter, a series of ten transition elements (lutetium through mercury) follows, and finally six main-group elements (thallium through radon) complete the period.

The seventh row is likewise analogous to the sixth row: 7s fills, then 5f, then 6d, and finally 7p, with a few anomalies along the way. For a very long time, the seventh row was incomplete as most of its elements do not occur in nature. The missing elements beyond uranium started to be synthesised in the laboratory in 1940, when neptunium was made. The row was completed with the synthesis of tennessine in 2010 (the last element oganesson had already been made in 2002), and the last elements in this seventh row were validated and given names in 2016.

This completes the modern periodic table, with all seven rows completely filled to capacity.

The following table shows the electron configuration of a neutral gas-phase atom of each element. Different configurations can be favoured in different chemical environments. The main-group elements have entirely regular electron configurations; the transition and inner transition elements show a number of irregularities due to the aforementioned competition between subshells close in energy level. For the last ten elements (109–118), experimental data is lacking and therefore calculated configurations have been shown instead. Completely filled subshells have been greyed out.

## External Links

The Periodic Table of Elements website on PubChem

Periodic table on Wikipedia

WebElements' official website, shows the periodic table at the main page.

Ptable, an interactive periodic table website.