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In 2019, Peter Atkins, an English chemist, claimed that “The periodic table is arguably the most important concept in chemistry” (Chemical Solutions Laboratories, 2019).  But why is it? Upon first glance, the periodic table does not look much like a conventional table; its shape and structure look jarring, and the complicated jumble of letters and numbers can be intimidating. Despite this, the periodic table is used by chemists to understand how atoms’ chemical properties relate to each other and how they would react. This is hugely important since chemical reactions are an essential part of chemistry itself, so the periodic table can never be missed out in any chemist’s lab. In this article, 3 main points about the periodic table will be discussed: its origin, the meaning of some of its features, and the trends that the elements follow. Hopefully, these will improve your understanding of how the periodic table came to be and why it’s so fundamental in chemistry. 


The origin of the periodic table

Although many scientists, such as John Newlands, attempted to sort the elements into tables, it was ultimately Dmitri Mendeleev who successfully created the basis for the periodic table in 1869. By ordering the elements in a row in order of atomic mass, he discovered that this row had recurring similar properties at certain periods. This became later known as the ‘periodic law’ (Russell, 2019). He also left gaps (seen by the question marks in the image below) in the periodic table as he predicted that new elements had to be discovered to fill in the gaps. This foresight proved to be correct when these gaps led to the discovery of elements like scandium and germanium (Brunning et al, 2022). 


At the time, noble gases were still not discovered (since they were discovered by William Ramsay in 1894), so there was no indication of them in Mendeleev’s periodic table. So when the noble gases were discovered and recognised, they created an 8th period instead of 7 which Mendeleev had originally put (Watson, 2021). 

However, Mendeleev’s periodic table had a problem. Assembling the elements in ascending order of atomic mass did not always make sense. For example, iodine has less mass than tellurium but ordering iodine before tellurium did not follow the existing pattern of chemical properties. Therefore, Mendeleev decided to switch elements when necessary. This reasoning was justified when the atomic number (the number of protons in the nucleus) was discovered by Henry Moseley. By firing a beam of electrons at different metals, a spectrum of X-rays could be detected, and its frequencies could indicate the positive charge of the nucleus, which was the number of protons. With this discovery made, Moseley claimed that the atomic number was the factor that decided an element’s characteristic - and he was right as this solved the problem with iodine and tellurium mentioned earlier. Ultimately, this enabled the periodic table to be arranged the way it is today: by ascending atomic number, not atomic mass (Brunning et al, 2022). 

Figure 1 Mendeleev’s periodic table

The features of the periodic table and what they tell you


Figure 2 the periodic table

As seen by the periodic table above, information is immediately given about the symbol, type, atomic number and relative atomic mass of the element. The left side of the periodic table (discluding hydrogen) and the middle are classified as metals, whilst the right side has elements classified as metalloids or nonmetals. These physical properties are easily observed, however, there is more that can be uncovered from the periodic table. Each period represents how many shells there are in an atom of the element and the group represents how many electrons are on the valence (outer) shell. For example, Mg (magnesium) would have 3 shells and 2 valence electrons since it is in the third period and second group of the periodic table. So, all group 1 metals would have the same number of electrons in their valence shell (1 electron) which is why all those elements have similar chemical properties as each other. Therefore, it can be said that elements in the same group have similar chemical properties. 


Figure 3 the periodic table with orbital shapes

In this version of the periodic table, certain ‘blocks’ have been colour coded to represent the orbital shapes. There are 4 orbital shapes in total in descending order of simplicity: s, p, d and f. Every orbital shape has at least 1 variation. The s orbital has 1; the p orbital has 3; the d orbital has 5; the f orbital has 7. These variations can be spotted in the gallery of atomic orbitals below.


Orbitals are important as they let scientists understand the distribution of electrons, and thus understand the electronic and optical properties of substances (Clean Energy Wiki, 2011).

Figure 4 the orbitron gallery of atomic orbitals

The periodic trends


Figure 5 the periodic trends


  • As you move across the period, electronegativity increases. This is because more protons (or a larger charge) in the nucleus means a higher attractive force between the bonded pair of electrons and protons, so it’s more likely to attract electrons towards itself.

  • As you move down a group, electronegativity decreases since the atom has a larger radius, and electron shielding reduces the electrostatic attraction between the atomic nucleus and valence electron(s).

Ionisation energy

  • As you move across the period, the ionisation energy increases because the electrons are harder to remove due to the increased electrostatic attraction between protons and electrons. This means more ionisation energy will be required to remove the electron away from the atom.

  • As you move down a group, ionisation energy decreases since the valence electron(s) is easier to remove from atoms that have more electron shielding. So less energy is required to do so.

Atomic radius

  • As you move across the period, the atomic radius decreases. This is because as atoms gain protons in the nucleus, the electrostatic attraction between the nucleus and valence electron(s) increases, which results in the electron(s) being pulled closer, so the radius decreases.

  • As you move down a group, the atomic radius increases. This is because the electrostatic attraction between the nucleus and valence electron(s) is weakened due to electron shielding as more shells are added.

Electron affinity

  • As you move across the period, the electron affinity increases since an increased number of protons in the nucleus mean that there is a higher electrostatic attraction between the nucleus and electrons, so the atom is more likely to gain electrons.

  • As you move down a group, the electron affinity decreases because the nucleus is further away from the valence shell, meaning that the electrostatic attraction between the nucleus and shell is weaker. This makes it harder for the nucleus to attract electrons towards itself to the valence shell, so the electron affinity decreases (University of Alabama in Huntsville, 2019).

Metallic and non-metallic character

As seen by the ionisation energy, electron affinity and electronegativity, metals have an easier time losing electrons and this gets progressive as we go down the group. Therefore, the metallic character is greatest at the bottom left of the periodic table, and that is where the most reactive metals are with francium being the most reactive. On the other hand, non-metals have a harder time losing electrons and going up the group only makes it harder. So, the most reactive nonmetals reside at the top right corner with fluorine being the most reactive. This excludes the noble gases as they have a full valence shell, meaning that they are inert (Helmenstine, 2022).


With Mendeleev’s prediction of the periodic law and Moseley’s discovery of the atomic number, the periodic table came to be. Without it, finding relationships between different elements and how they would react with each other would be significantly more difficult, and useful trends would have never been recognised. Therefore, Peter Atkins was correct. The periodic table can be argued to be the most important concept in chemistry. So next time you use the periodic table in your chemistry lesson (or any science lesson), just know that you are using one of the most important tools a chemist could ever own. 



Electronegativity: A chemical property that describes the tendency for an atom or a functional group to attract a bonding pair of electrons (Clark, 2019).

Ionisation energy: Energy required to remove an electron away from an atom in the gaseous state.

Atomic radius: Half the distance between the nuclei of two atoms of the same element.

Electron affinity: How able the atom is to accept an electron (University of Alabama in Huntsville, 2019).

Metallic character: How reactive a metal is, or how well a metal can lose electrons to create a full valence shell.

Non-metallic character: How reactive a non-metal is, or how well a non-metal can gain electrons to create a full valence shell (Helmenstine, 2019).

Electron shielding: When the inner electrons block the valence electron(s) from the nucleus.


Written By Ari Kim (12W), a student at Bangkok Patana School. 

Edited by: Samuel Lim (13L)




Chemical Solutions Laboratories, Inc. (2019), Happy 150th Birthday to the Periodic Table! [online] Available at:


Russell J. (2019), Elementary: The Periodic Table Explained, Michael O’Mara Books Limited [book]


Brunning A. et al (2022), The Chemistry Book: Big Ideas Simply Explained, D.K. London: Penguin Random House


Watson K. (September 2021), Sir William Ramsay. [online] Available at:


Clean Energy Wiki (May 2011), Atomic Orbitals and Nodes. [online] Available at:


University of Alabama in Huntsville (2019), periodic_trends. [online] Available at:


Helmenstine A. (May 2022), Metallic Character Trend on the Periodic Table. [online] Available at:


Clark J. (November 2019), Electronegativity. [online] Available at:


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