Some elements exist in several different structural forms, called allotropes. Each allotrope has different physical properties.

For more information on the Visual Elements image see the Uses and properties section below.



A vertical column in the periodic table. Members of a group typically have similar properties and electron configurations in their outer shell.

A horizontal row in the periodic table. The atomic number of each element increases by one, reading from left to right.

Elements are organised into blocks by the orbital type in which the outer electrons are found. These blocks are named for the characteristic spectra they produce: sharp (s), principal (p), diffuse (d), and fundamental (f).

Atomic number
The number of protons in an atom.

Electron configuration
The arrangements of electrons above the last (closed shell) noble gas.

Melting point
The temperature at which the solid–liquid phase change occurs.

Boiling point
The temperature at which the liquid–gas phase change occurs.

The transition of a substance directly from the solid to the gas phase without passing through a liquid phase.

Density (g cm−3)
Density is the mass of a substance that would fill 1 cm3 at room temperature.

Relative atomic mass
The mass of an atom relative to that of carbon-12. This is approximately the sum of the number of protons and neutrons in the nucleus. Where more than one isotope exists, the value given is the abundance weighted average.

Atoms of the same element with different numbers of neutrons.

CAS number
The Chemical Abstracts Service registry number is a unique identifier of a particular chemical, designed to prevent confusion arising from different languages and naming systems.

Fact box

Group 18  Melting point Unknown 
Period Boiling point −268.928°C, −452.07°F, 4.222 K 
Block Density (g cm−3) 0.000164 
Atomic number Relative atomic mass 4.003  
State at 20°C Gas  Key isotopes 4He 
Electron configuration 1s2  CAS number 7440-59-7 
ChemSpider ID 22423 ChemSpider is a free chemical structure database


Image explanation

Murray Robertson is the artist behind the images which make up Visual Elements. This is where the artist explains his interpretation of the element and the science behind the picture.


The description of the element in its natural form.

Biological role

The role of the element in humans, animals and plants.

Natural abundance

Where the element is most commonly found in nature, and how it is sourced commercially.

Uses and properties

Image explanation
The image is of the sun because helium gets its name from ‘helios’, the Greek word for the sun. Helium was detected in the sun by its spectral lines many years before it was found on Earth.
A colourless, odourless gas that is totally unreactive.
Helium is used as a cooling medium for the Large Hadron Collider (LHC), and the superconducting magnets in MRI scanners and NMR spectrometers. It is also used to keep satellite instruments cool and was used to cool the liquid oxygen and hydrogen that powered the Apollo space vehicles.

Because of its low density helium is often used to fill decorative balloons, weather balloons and airships. Hydrogen was once used to fill balloons but it is dangerously reactive.

Because it is very unreactive, helium is used to provide an inert protective atmosphere for making fibre optics and semiconductors, and for arc welding. Helium is also used to detect leaks, such as in car air-conditioning systems, and because it diffuses quickly it is used to inflate car airbags after impact.

A mixture of 80% helium and 20% oxygen is used as an artificial atmosphere for deep-sea divers and others working under pressurised conditions.

Helium-neon gas lasers are used to scan barcodes in supermarket checkouts. A new use for helium is a helium-ion microscope that gives better image resolution than a scanning electron microscope.
Biological role
Helium has no known biological role. It is non-toxic.
Natural abundance
After hydrogen, helium is the second most abundant element in the universe. It is present in all stars. It was, and is still being, formed from alpha-particle decay of radioactive elements in the Earth. Some of the helium formed escapes into the atmosphere, which contains about 5 parts per million by volume. This is a dynamic balance, with the low-density helium continually escaping to outer space.

It is uneconomical to extract helium from the air. The major source is natural gas, which can contain up to 7% helium.
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In 1868, Pierre J. C. Janssen travelled to India to measure the solar spectrum during a total eclipse and observed a new yellow line which indicated a new element. Joseph Norman Lockyer recorded the same line by observing the sun through London smog and, assuming the new element to be a metal, he named it helium.

In 1882, the Italian Luigi Palmieri found the same line the spectrum of gases emitted by Vesuvius, as did the American William Hillebrand in 1889 when he collected the gas given off by the mineral uraninite (UO2) as it dissolves in acid. However, it was Per Teodor Cleve and Nils Abraham Langer at Uppsala, Sweden, in 1895, who repeated that experiment and confirmed it was helium and measured its atomic weight.

Atomic radius, non-bonded
Half of the distance between two unbonded atoms of the same element when the electrostatic forces are balanced. These values were determined using several different methods.

Covalent radius
Half of the distance between two atoms within a single covalent bond. Values are given for typical oxidation number and coordination.

Electron affinity
The energy released when an electron is added to the neutral atom and a negative ion is formed.

Electronegativity (Pauling scale)
The tendency of an atom to attract electrons towards itself, expressed on a relative scale.

First ionisation energy
The minimum energy required to remove an electron from a neutral atom in its ground state.

Atomic data

Atomic radius, non-bonded (Å) 1.400 Covalent radius (Å) 0.37
Electron affinity (kJ mol−1) Not stable Electronegativity
(Pauling scale)
Ionisation energies
(kJ mol−1)


Common oxidation states

The oxidation state of an atom is a measure of the degree of oxidation of an atom. It is defined as being the charge that an atom would have if all bonds were ionic. Uncombined elements have an oxidation state of 0. The sum of the oxidation states within a compound or ion must equal the overall charge.


Atoms of the same element with different numbers of neutrons.

Key for isotopes

Half life
  y years
  d days
  h hours
  m minutes
  s seconds
Mode of decay
  α alpha particle emission
  β negative beta (electron) emission
  β+ positron emission
  EC orbital electron capture
  sf spontaneous fission
  ββ double beta emission
  ECEC double orbital electron capture

Oxidation states and isotopes

Common oxidation states
Isotopes Isotope Atomic mass Natural abundance (%) Half life Mode of decay
  3He 3.016 0.000134
  4He 4.003 99.9999


Data for this section been provided by the British Geological Survey.

Relative supply risk

An integrated supply risk index from 1 (very low risk) to 10 (very high risk). This is calculated by combining the scores for crustal abundance, reserve distribution, production concentration, substitutability, recycling rate and political stability scores.

Crustal abundance (ppm)

The number of atoms of the element per 1 million atoms of the Earth’s crust.

Recycling rate

The percentage of a commodity which is recycled. A higher recycling rate may reduce risk to supply.


The availability of suitable substitutes for a given commodity.
High = substitution not possible or very difficult.
Medium = substitution is possible but there may be an economic and/or performance impact
Low = substitution is possible with little or no economic and/or performance impact

Production concentration

The percentage of an element produced in the top producing country. The higher the value, the larger risk there is to supply.

Reserve distribution

The percentage of the world reserves located in the country with the largest reserves. The higher the value, the larger risk there is to supply.

Political stability of top producer

A percentile rank for the political stability of the top producing country, derived from World Bank governance indicators.

Political stability of top reserve holder

A percentile rank for the political stability of the country with the largest reserves, derived from World Bank governance indicators.

Supply risk

Relative supply risk 6.5
Crustal abundance (ppm) 0.008
Recycling rate (%) Unknown
Substitutability Unknown
Production concentration (%) 22.2
Reserve distribution (%) 21
Top 3 producers
  • 1) USA
  • 2) Algeria
  • 3) Russia
Top 3 reserve holders
  • 1) USA
  • 2) Qatar
  • 3) Algeria
Political stability of top producer 56.6
Political stability of top reserve holder 56.6


Specific heat capacity (J kg−1 K−1)

Specific heat capacity is the amount of energy needed to change the temperature of a kilogram of a substance by 1 K.

Young's modulus

A measure of the stiffness of a substance. It provides a measure of how difficult it is to extend a material, with a value given by the ratio of tensile strength to tensile strain.

Shear modulus

A measure of how difficult it is to deform a material. It is given by the ratio of the shear stress to the shear strain.

Bulk modulus

A measure of how difficult it is to compress a substance. It is given by the ratio of the pressure on a body to the fractional decrease in volume.

Vapour pressure

A measure of the propensity of a substance to evaporate. It is defined as the equilibrium pressure exerted by the gas produced above a substance in a closed system.

Pressure and temperature data – advanced

Specific heat capacity
(J kg−1 K−1)
5193 Young's modulus (GPa) Unknown
Shear modulus (GPa) Unknown Bulk modulus (GPa) Unknown
Vapour pressure  
Temperature (K)
400 600 800 1000 1200 1400 1600 1800 2000 2200 2400
Pressure (Pa)
- - - - - - - - - - -
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Listen to Helium Podcast
Transcript :

Chemistry in its element: helium


You're listening to Chemistry in its element brought to you by Chemistry World, the magazine of the Royal Society of Chemistry.

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Chris Smith

Hello, this week we're almost at the top of the periodic table because we're taking a look at the lighter than air gas helium. But for this chemist a helium filled bobbing balloon is actually a source of pain and not a source of pleasure. Here's Peter Wothers.

Peter Wothers

We are all familiar with the lighter-than-air gas helium, but whenever I see a balloon floating on a string, I feel a little sad. It's not because I'm a miserable old so-and-so - it's just because, unlike the happy child on the other end of the string, I am aware of the valuable resource that's about to be lost forever.

Helium is the second most abundant element in the universe, but here on earth, it's rather rare. Most people guess that we extract helium from the air, but actually we dig it out of the ground. Helium can be found in certain parts of the world, notably in Texas, as a minor component in some sources of natural gas. The interesting thing is how this gas gets into the ground in the first place. Unlike virtually every other atom around us, each atom of helium has been individually formed after the formation of the earth.

The helium is formed during the natural radioactive decay of elements such as uranium and thorium. These heavy elements were formed before the earth but they are not stable and very slowly, they decay. One mode of decay for uranium is to emit an alpha-particle. This alpha-particle is actually just the heart of a helium atom - its nucleus. Once it has grabbed a couple of electrons, a helium atom has been born.

This decay process for uranium is incredibly slow; the time it takes a given quantity of uranium to halve, its so-called half-life, is comparable to the age of the earth. This means that helium has been continuously generated ever since the earth was formed. Some of the gas might eventually creep through the earth and escape into the atmosphere; fortunately, when conditions are right, some is trapped underground and can be harvested for our use.

The situation is very different in space. The sun is comprised of about 75% by mass of hydrogen and 24% of helium. The remaining one percent is made up of all the heavier elements. In the high temperatures of the sun, the hydrogen nuclei are fused together to eventually form helium. This fusion process, whereby heavier atoms are made from lighter ones, liberates vast amounts of energy. Recreating the process on earth may be the answer to our energy problems in the future.

Since helium makes up about a quarter of the mass of the sun, it is not surprising that its presence was detected there over 100 years ago. What is perhaps surprising, is that helium was discovered in space 26 years before it was found on earth.

It has been known for hundreds of years that certain elements impart characteristic colours to a flame - a fact crucial to the coloured fireworks that we enjoy. Copper, for example, gives a green colour, whereas sodium gives a yellow colour. It is actually possible to identify elements by the careful examination of such coloured flames. The light is split up into a spectrum using a prism or diffraction grating in an instrument called a spectroscope. Rather than seeing a continuous rainbow of colours, a series of sharp coloured lines is formed. This series of lines is characteristic of the particular element and acts as a sort of fingerprint.

In the 19th century, scientists turned their spectroscopes to the sun and began to detect certain metals there, including sodium, magnesium, calcium and iron. In 1868 two astronomers, Janssen and Lockyer, independently noticed some very clear lines in the solar spectrum that did not match up to any known metals. While other astronomers of the time were unsure, Lockyer suggested these unidentified lines belonged to a new metal which he named Helium after the Greek personification of the sun, Helios. For over 20 years, no sign of the metal helium was detected on earth and Lockyer began to be mocked for his mythical element. However, in 1895 the chemist William Ramsay detected helium in the gas given out when a radioactive mineral of uranium was treated with acid. The helium formed from the radioactive decay had been trapped in the rock but liberated when the rock was dissolved away in the acid.

Finally Lockyer's element had been discovered on earth, but it was no metal, rather an extremely unreactive gas. To this day, helium remains the only non-metal whose name ends with the suffix -ium, an ending otherwise exclusively reserved for metals.

Aside from being used to fill balloons, both for our entertainment, and for more serious purposes, such as for weather balloons, helium is used in other applications which depend on its unique properties. Being so light, and yet totally chemically inert, helium can be mixed with oxygen in order to make breathing easier. This mixture, known as heliox, can help save new-born babies with breathing problems, or help underwater divers safely reach the depths of the oceans. At minus 269 degrees centigrade, liquid helium has the lowest boiling point of any substance. Because of this, it is used to provide the low temperatures needed for superconducting magnets, such as those used in most MRI scanners in hospitals.

In many facilities where helium is used, it is captured and reused. If it isn't, it escapes into the air. But it doesn't simply accumulate in the atmosphere. Helium is so light that it can escape the pull of the earth's gravitational field and leave our planet forever. This is the fate of the helium in our balloons. Whereas it may be possible to reclaim and recycle other elements that we have used and discarded, when we waste helium, it is lost for good. In 100 years time, people will look back with disbelief that we wasted this precious, unique element by filling up party balloons.

Chris Smith

Cambridge University's Peter Wothers telling us the tale of element number two, Helium. Next time we're off to 18th century Scotland and an element that was the wrong colour.

Richard Van Noorden

In 1787, an intriguing mineral came to Edinburgh from a Lead mine in a small village on the shores of Loch Sunart, Argyll. At that time, the stuff was thought to be some sort of Barium compound. Other chemists, such as Edinburgh's Thomas Hope later prepared a number of compounds with the element, noting that it caused the candle's flame to burn red, while Barium compounds gave a green colour.

Chris Smith

And that's because it wasn't Barium at all, it was Strontium and Richard Van Noorden will be here to explain how, amongst other things, it's shown us that Roman gladiators weren't meat eaters they were in fact vegetarians. That's next week's Chemistry in its Element and I hope you can join us. I'm Chris Smith, thank you for listening and goodbye.


Chemistry in its element is brought to you by the Royal Society of Chemistry and produced by There's more information and other episodes of Chemistry in its element on our website at

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Visual Elements images and videos
© Murray Robertson 2011.



W. M. Haynes, ed., CRC Handbook of Chemistry and Physics, CRC Press/Taylor and Francis, Boca Raton, FL, 95th Edition, Internet Version 2015, accessed December 2014.
Tables of Physical & Chemical Constants, Kaye & Laby Online, 16th edition, 1995. Version 1.0 (2005), accessed December 2014.
J. S. Coursey, D. J. Schwab, J.J. Tsai, and R. A. Dragoset, Atomic Weights and Isotopic Compositions(version 3.0), 2010, National Institute of Standards and Technology, Gaithersburg, MD, accessed December 2014.
T. L. Cottrell, The Strengths of Chemical Bonds, Butterworth, London, 1954.


Uses and properties

John Emsley, Nature’s Building Blocks: An A-Z Guide to the Elements, Oxford University Press, New York, 2nd Edition, 2011.
Thomas Jefferson National Accelerator Facility - Office of Science Education, It’s Elemental - The Periodic Table of Elements, accessed December 2014.
Periodic Table of Videos, accessed December 2014.


Supply risk data

Derived in part from material provided by the British Geological Survey © NERC.


History text

© John Emsley 2012.



Produced by The Naked Scientists.


Periodic Table of Videos

Created by video journalist Brady Haran working with chemists at The University of Nottingham.
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