Periodic Table of Elements

How to read the periodic table

Every tile packs four facts into a square. The small number in the corner is the atomic number — the count of protons in the nucleus, which is what makes an element that element. Change it and you have a different element entirely. The large letters are the chemical symbol, an international shorthand that is the same in every language, though several are abbreviations of Latin names rather than English ones: Fe for iron comes from ferrum, Au for gold from aurum, Pb for lead from plumbum. Below that sits the element name, and below that the atomic mass in unified atomic mass units.

The layout itself carries information. Read across a row and the atomic number climbs by one at each step while properties shift gradually — metals on the left give way to non-metals on the right. Read down a column and you find elements that behave alike, because they have the same number of electrons in their outermost shell. That repetition, appearing again and again as atomic number increases, is the periodicity the table is named for.

How the elements are arranged

The table has three organising ideas layered on top of each other.

• Periods are the seven horizontal rows. The period number tells you how many electron shells the atom has: everything in period 3 has three shells. Rows get longer as you go down because bigger shells hold more electrons — 2, then 8, then 8, then 18.
• Groups are the eighteen vertical columns. Elements in a group share their outer-electron count and therefore their chemistry. Several have names people actually use: group 1 are the alkali metals, group 17 the halogens, group 18 the noble gases.
• Blocks are the regions named after the orbital being filled — s, p, d and f. The s-block is the two columns on the left, the p-block the six on the right, the d-block the ten transition-metal columns in the middle, and the f-block the two rows pulled out below.

That last one explains the table's odd shape. The f-block genuinely belongs between groups 2 and 3 in periods 6 and 7, but putting it there makes the table 32 columns wide — too wide for a textbook page or a classroom wall. So it is conventionally cut out and parked underneath, with a “57–71” and “89–103” placeholder left behind. Switch this table to its 32-column view above to put it back where it belongs.

The types of element

Elements are grouped into families that share a way of behaving. These categories are a convention rather than a law of nature — chemists disagree at the edges, particularly over which elements count as metalloids — but they capture real patterns.

What each property means

Use the Colour by control above to paint the whole table with any of these and watch the trend appear.

Atomic mass

The weighted average mass of the element's naturally occurring isotopes, in unified atomic mass units (u). It is not a whole number because most elements are a mixture of isotopes. For elements with no stable isotope, the table shows the mass number of the longest-lived one in square brackets instead.

Electronegativity

How hard an atom pulls on shared electrons in a bond, on the Pauling scale. It rises as you go right and falls as you go down, peaking at fluorine (3.98) — the greediest element there is. The difference in electronegativity between two atoms predicts whether their bond will be covalent, polar or ionic.

Ionisation energy

The energy needed to tear the outermost electron away. Noble gases sit at the peaks because their full shells are stable; alkali metals sit in the troughs because their single outer electron is barely held.

Atomic radius

How big the atom is. Counter-intuitively, atoms get smaller across a period: the extra protons pull the same shell in tighter. They jump larger at the start of each new row, when a fresh shell opens.

Oxidation states

The charges the element takes in compounds — see the full charges table. Main-group elements are predictable from their group number; transition metals have several, which is what makes their compounds so colourful.

A short history

By the 1860s chemists knew of about 60 elements and could see that something ordered them, but nobody had made the pattern stick. Dmitri Mendeleev published his version in 1869, and what made it work was a piece of nerve: where the pattern demanded an element that nobody had ever seen, he left a gap and predicted what would fill it — its mass, its density, the colour of its oxide. Gallium arrived in 1875, scandium in 1879 and germanium in 1886, each close to his predictions. The gaps had made the table testable, and it passed.

Mendeleev ordered by atomic weight, which mostly worked but produced a few stubborn reversals. Henry Moseley resolved them in 1913 by showing that the real ordering principle is nuclear charge — the atomic number — not weight. Everything since has been filling in: the noble gases in the 1890s, which needed a whole new column; the synthetic elements from 1937 onward; and finally elements 113, 115, 117 and 118, named in 2016, which completed period 7.