The Periodic Table
One chart that sorts every known element and hints at how each one behaves
Every material you have ever touched, from your phone screen to the air in your lungs, is built from just 118 kinds of atoms. The periodic table lists all of them on one page. It is not just a list, though. Where an element sits tells you a surprising amount about it: whether it is a shiny metal or a brittle powder, whether it reacts violently with water or ignores almost everything, and how big its atoms are.
On this page we'll cover the atom basics you need to read the table, how a Russian chemist named Dmitri Mendeleev pieced it together, how its rows and columns are organized, and the patterns, called periodic trends, that run across it. Then you can explore a clickable table of all 118 elements and try the quiz.
Reading an Element Tile
Just enough atom to make sense of the boxes.
An atom has a tiny, heavy nucleus made of protons (charge +1) and neutrons (no charge). Around it move electrons (charge −1), which are about 1,836 times lighter than a proton. Protons and neutrons each have a mass of about 1 atomic mass unit (u).
Atomic number is the number of protons. It is the element’s ID: every atom with 6 protons is carbon, every atom with 79 is gold. A neutral atom has as many electrons as protons.
Mass number is protons plus neutrons for one particular atom. Atoms of the same element can have different numbers of neutrons. These versions are called isotopes. Carbon-12 has 6 neutrons, and carbon-14 has 8.
The atomic mass printed on the table is an average over the isotopes found in nature. Chlorine shows 35.45 because about 76% of chlorine atoms are chlorine-35[11] and about 24% are chlorine-37.
Electrons sit in shells (energy levels) around the nucleus. For the first 20 elements you can fill them simply: 2 electrons in the first shell, then 8, then 8, and then the fourth shell begins. Sodium’s 11 electrons go 2, 8, 1.
The electrons in the outermost shell are the valence electrons. They are the ones that react, so they decide most of an element’s chemistry. Here is the payoff: for the main-group elements, the column tells you the count. Group 1 elements have 1 valence electron, group 2 have 2, and groups 13 to 18 have 3 to 8 (helium, with only 2, is the exception).
Real electrons don’t travel in neat circles. They live in fuzzy regions called orbitals, which you’ll meet later. The shell picture is still a great tool for reading the table.
Mendeleev’s Big Idea
How a chemist writing a textbook ended up predicting elements no one had seen.
By the 1860s chemists knew 63 elements and could see family resemblances. Several people, including John Newlands in England, noticed that properties seemed to repeat when elements were lined up by atomic weight.[2]
Dmitri Mendeleev, a professor in St. Petersburg, was writing a chemistry textbook and wanted a sensible order for the chapters. He arranged the elements by atomic weight in a grid, so that elements with similar properties lined up together. His table was presented to the Russian Chemical Society in March 1869.[1] (The German chemist Lothar Meyer built a very similar table around the same time.)
Mendeleev’s bold move was trusting the pattern over the list. When an element didn’t fit, he left a gap and said an undiscovered element belonged there. Within two years he had predicted the properties of three missing elements in detail. Below aluminum he described “eka-aluminium”, and in 1875 gallium was discovered with almost exactly the density and behavior he forecast. “Eka-silicon” turned up as germanium in 1886, with an atomic weight of about 72.6, close to his prediction of 72. Scandium (1879) filled a third gap.[1][3]
He also swapped a few elements out of weight order when their properties demanded it, such as putting tellurium (127.6) before iodine (126.9). In 1913 Henry Moseley used X-rays to show why that worked: the true order is by atomic number, the number of protons, not by weight. That is how the modern table is arranged.[5]
The table kept growing. A whole column, the noble gases, was added after argon was found in 1894.[7] Glenn Seaborg, who helped make many of the elements past uranium, proposed giving the actinides their own row.[8] The seventh row was finally completed when the names nihonium, moscovium, tennessine and oganesson were approved in 2016.[10] Element 101 is called mendelevium in Mendeleev’s honor.
Groups, Periods and Blocks
Columns are families, rows are shells.
Groups (columns)
There are 18 groups, numbered 1 to 18 from left to right. Elements in a group have the same number of valence electrons, so they react in similar ways. Sodium and potassium both lose one electron easily; chlorine and bromine both grab one.
Periods (rows)
There are 7 periods. The period number tells you how many electron shells an atom has. Hydrogen and helium (period 1) have one shell; sodium through argon (period 3) have three. Moving left to right along a period, each element has one more proton and one more electron than the last.
Blocks
The table’s odd shape comes from which type of orbital is being filled. The s-block is groups 1 and 2 (plus helium), the p-block is groups 13 to 18, the d-block is groups 3 to 12 (the transition metals), and the f-block is the two rows usually printed underneath. Try “Color by block” in the interactive table.
Why two rows float at the bottom
The lanthanides and actinides really belong inside periods 6 and 7. Squeezing them in would make the table 32 columns wide, so they are pulled out and printed below to keep it page-shaped. Chemists still debate exactly which elements sit in group 3 (lanthanum and actinium, or lutetium and lawrencium). This page shows the full rows of 15 below the table.
Metals, Nonmetals and Metalloids
A staircase line splits the table into two very different worlds.
A zig-zag “staircase” runs down the right side of the table from boron to astatine. About three-quarters of the elements, everything to the left of it (except hydrogen), are metals. Nonmetals sit to the upper right, and the elements hugging the line are metalloids.
| Metals | Nonmetals | Metalloids | |
|---|---|---|---|
| Look | Shiny (lustrous) | Dull, many are gases | Often shiny but brittle |
| Conduct heat and electricity? | Yes, well | Poorly (graphite is an exception) | Somewhat: semiconductors |
| Bend or shatter? | Malleable and ductile | Brittle as solids | Brittle |
| With electrons | Tend to lose them, forming + ions | Tend to gain or share them | In between |
| Examples | Iron, copper, sodium, gold | Oxygen, carbon, sulfur, neon | Boron, silicon, germanium, arsenic, antimony, tellurium |
Silicon’s in-between ability to conduct is why it sits inside nearly every computer chip. At room temperature only two elements are liquids, mercury (a metal) and bromine (a nonmetal), and eleven are gases.
Element Families
Some columns are close enough relatives to get their own names.
Alkali metals (group 1)
Lithium, sodium, potassium, rubidium, cesium, francium. Soft enough to cut with a knife and so reactive they are stored under oil. One valence electron, which they give away eagerly, so they react with water to make hydrogen gas and a strong base. Hydrogen sits above them but is a nonmetal.
Alkaline earth metals (group 2)
Beryllium to radium. Two valence electrons, reactive but less wild than group 1. Calcium builds your bones and teeth, and magnesium burns with a dazzling white light.
Transition metals (groups 3 to 12)
The classic “metal” metals: iron, copper, nickel, silver, gold, titanium. Mostly hard, strong and good conductors. Many form colored compounds; copper sulfate is bright blue.
Halogens (group 17)
Fluorine, chlorine, bromine, iodine, astatine (and tennessine). Seven valence electrons, so they are one short of a full shell and grab electrons from metals to make salts. The name means “salt-former”: sodium + chlorine gives table salt, NaCl.
Noble gases (group 18)
Helium, neon, argon, krypton, xenon, radon. Their outer shells are already full, so they barely react at all. That’s why helium is safe in balloons and argon is used to protect hot metal during welding.
Lanthanides and actinides
The two bottom rows. Lanthanides (often called rare earths) go into strong magnets, phone screens and wind turbines. Every actinide is radioactive; uranium and plutonium fuel nuclear reactors, and most of the later ones have only ever been made in labs.
Periodic Trends
The patterns that make the table more than a list.
Almost every trend comes down to a tug-of-war. The nucleus pulls on the outer electrons. Moving across a period, protons are added but the electrons stay in the same shell, so the pull gets stronger. Moving down a group, each new shell puts the outer electrons farther away, and the inner electrons shield them from the nucleus, so the pull gets weaker.
Atomic radius
The size of an atom. It decreases across a period (the stronger pull draws the electrons in) and increases down a group (more shells). A cesium atom is several times wider than a fluorine atom.
Ionization energy
The energy needed to pull one electron off an atom. It increases across a period and decreases down a group. Sodium gives up its electron for about 496 kJ/mol, chlorine needs about 1,251, and helium holds the record at about 2,372.[12]
Electronegativity
How strongly an atom attracts the shared electrons in a chemical bond. It increases across a period and decreases down a group. Fluorine is the most electronegative element (3.98 on the Pauling scale)[13]; cesium and francium are the least. A few heavy elements break the “decreases down a group” pattern: tin is 1.96 but lead is 2.33, tungsten (2.36) beats molybdenum (2.16), and gold (2.54) beats silver (1.93). Color the table by electronegativity to see it.
Reactivity
This one flips depending on the side. Metals react by losing electrons, so they get more reactive going down: lithium fizzes in water, sodium skates and pops, potassium bursts into a lilac flame, and cesium explodes. Nonmetals react by gaining electrons, so they get more reactive going up: fluorine is the most reactive halogen, iodine the least of the common ones.
How Do We Know?
You can’t see an atom. So how do scientists tell one element from another, or prove a new one exists?
Every element has a light “barcode”
Heat an element until it glows and pass the light through a prism. Instead of a full rainbow you get a few sharp, bright emission lines. Each line is light given off when an electron drops from a higher energy level to a lower one. Every element has its own pattern, like a fingerprint. Studying these lines is called spectroscopy.
Original diagram. Line positions are rounded to the nearest nanometre, from standard wavelength tables.
In 1860 and 1861 Robert Bunsen and Gustav Kirchhoff used their new spectroscope to discover two metals, cesium (1860) and rubidium (1861), from lines no known element made.[4] In 1868 astronomers saw an unknown yellow line in sunlight. It belonged to helium, named after the Greek word for the Sun, which was not found on Earth until years later.[6]
X-rays reveal the atomic number
In 1913–14 Henry Moseley fired electrons at samples of many elements and measured the X-rays they gave off. The X-ray frequency rose in neat steps from one element to the next, matching a whole number: the charge of the nucleus, which we now call the atomic number. He published the full series in two papers in 1913 and 1914. His results left gaps at 43, 61, 72 and 75, so he knew four elements were still missing. All were later found or made: hafnium (72) in 1923, rhenium in 1925, technetium in 1937 and promethium in 1945.[5][14]
Confirming a brand-new superheavy element
The heaviest elements don’t occur in nature at all. Labs such as Dubna in Russia and RIKEN in Japan make them by firing one kind of nucleus at another, often for weeks. Only a handful of atoms form, and each lasts a tiny fraction of a second. Researchers identify them by their decay chain: the new atom breaks down step by step, and the steps are traced to nuclei that are already known.
A claim is not accepted until an independent panel of chemists and physicists from IUPAC and IUPAP checks the evidence. In 2015 that panel confirmed elements 113, 115, 117 and 118, which completed the seventh row.[9][10]
The Interactive Periodic Table
Click or tab to any element to see its details. Use the menu to recolor the table, or a legend button to spotlight one family.
On a small screen, swipe sideways to see the whole table.
Atomic masses are IUPAC standard values, rounded. A number in [brackets] is the mass number of a long-lived isotope for elements with no stable isotopes. Electronegativities are Pauling values.[13] Sources disagree on a few borderline elements: some list polonium and astatine as metalloids. For elements past about 103, only a few atoms have ever been made, so their categories and states are predictions.
Test Yourself
Ten quick questions. Click an answer to check it. The table above is allowed!
The Lab Bench Challenge
Think About It
Questions to puzzle over in class or at home.
- Mendeleev left gaps instead of forcing every element into place. Why was that a braver, more scientific choice than making the list “complete”?
- Find sodium and chlorine on the table. Using valence electrons, explain why they combine so eagerly to make salt.
- Helium and neon are both noble gases, but helium has 2 valence electrons and neon has 8. Why do they still belong in the same group?
- If a brand-new element 119 were made, which group would it join, and what would you predict about how it reacts with water?
Sources
- Science History Institute, “An Element of Order” (Mendeleev’s 1869 table and his predictions). https://sciencehistory.org/stories/magazine/an-element-of-order/
- Philip Stewart, “Mendeleev’s predictions: success and failure,” Foundations of Chemistry (2019). https://link.springer.com/article/10.1007/s10698-018-9312-0
- Purdue University, “Dmitri Mendeleev” (eka-silicon predictions vs. germanium). https://chemed.chem.purdue.edu/genchem/history/mendeleeff.html
- Science History Institute, “Robert Bunsen and Gustav Kirchhoff.” https://www.sciencehistory.org/education/scientific-biographies/robert-bunsen-and-gustav-kirchhoff/
- University of Oxford Department of Physics, “Moseley and X rays.” https://www.physics.ox.ac.uk/about-us/our-history/highlights-archive/moseley-and-x-rays
- Science Museum Group, “150 years since the discovery of helium.” https://blog.sciencemuseum.org.uk/150-years-since-the-discovery-of-helium/
- Nobel Prize, “William Ramsay: Facts” (Chemistry 1904, the noble gases). https://www.nobelprize.org/prizes/chemistry/1904/ramsay/facts/
- Science History Institute, “Glenn Theodore Seaborg.” https://www.sciencehistory.org/education/scientific-biographies/glenn-theodore-seaborg/
- IUPAC, “Discovery and Assignment of Elements with Atomic Numbers 113, 115, 117 and 118” (2015). https://iupac.org/discovery-and-assignment-of-elements-with-atomic-numbers-113-115-117-and-118/
- IUPAC, “IUPAC is naming the four new elements nihonium, moscovium, tennessine, and oganesson” (2016). https://iupac.org/iupac-is-naming-the-four-new-elements-nihonium-moscovium-tennessine-and-oganesson/
- CIAAW, “Atomic weight of chlorine” (isotope abundances). https://www.ciaaw.org/chlorine.htm
- NIST Atomic Spectra Database, ionization energies. https://physics.nist.gov/PhysRefData/ASD/ionEnergy.html
- PubChem, “Periodic Table of Elements” (masses, electronegativity, states, discovery years). https://pubchem.ncbi.nlm.nih.gov/periodic-table/
- “Henry Moseley, X-ray spectroscopy and the periodic table,” Philosophical Transactions of the Royal Society A (2020). https://royalsocietypublishing.org/doi/10.1098/rsta.2019.0302
- U.S. EPA, “Protect Your Family from Sources of Lead” (1978 ban on lead paint in homes). https://www.epa.gov/lead/protect-your-family-sources-lead
- U.S. EPA, “EPA Takes Final Step in Phaseout of Leaded Gasoline” (1996). https://www.epa.gov/archive/epa/aboutepa/epa-takes-final-step-phaseout-leaded-gasoline.html