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 Melting point 696°C, 1285°F, 969 K 
Period Boiling point 1500°C, 2732°F, 1773 K 
Block Density (g cm−3)
Atomic number 88  Relative atomic mass [226]  
State at 20°C Solid  Key isotopes 226Ra 
Electron configuration [Rn] 7s2  CAS number 7440-14-4 
ChemSpider ID 4886483 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 represents the former use of radium in luminous paint used for clock and watch dials.
A soft, shiny and silvery radioactive metal.
Radium now has few uses, because it is so highly radioactive.

Radium-223 is sometimes used to treat prostate cancer that has spread to the bones. Because bones contain calcium and radium is in the same group as calcium, it can be used to target cancerous bone cells. It gives off alpha particles that can kill the cancerous cells.

Radium used to be used in luminous paints, for example in clock and watch dials. Although the alpha rays could not pass through the glass or metal of the watch casing, it is now considered to be too hazardous to be used in this way.
Biological role
Radium has no known biological role. It is toxic due to its radioactivity.
Natural abundance
Radium is present in all uranium ores, and could be extracted as a by-product of uranium refining. Uranium ores from DR Congo and Canada are richest in radium. Today radium is extracted from spent fuel rods from nuclear reactors. Annual production of this element is fewer than 100 grams per year.
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Radium was discovered in 1898 by Marie Curie and Pierre Curie. They managed to extract 1 mg of radium from ten tonnes of the uranium ore pitchblende (uranium oxide, U3O8), a considerable feat, given the chemically methods of separation available to them. They identified that it was a new element because its atomic spectrum revealed new lines. Their samples glowed with a faint blue light in the dark, caused by the intense radioactivity exciting the surrounding air.

The metal itself was isolated by Marie Curie and André Debierne in 1911, by means of the electrolysis of radium chloride. At Debierne’s suggestion, they used a mercury cathode in which the liberated radium dissolved. This was then heated to distil off the mercury leaving the radium behind.

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) 2.83 Covalent radius (Å) 2.11
Electron affinity (kJ mol−1) 9.65 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 2
Isotopes Isotope Atomic mass Natural abundance (%) Half life Mode of decay
  223Ra 223.019 - 11.43 d  α 
  224Ra 224.020 - 3.66 d  α 
  226Ra 226.025 - 1599 y  α 
        > 4 x 1018 sf 
  228Ra 228.031 - 5.76 y  β- 


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 Unknown
Crustal abundance (ppm) 0.0000009
Recycling rate (%) Unknown
Substitutability Unknown
Production concentration (%) Unknown
Reserve distribution (%) Unknown
Top 3 producers
  • Unknown
Top 3 reserve holders
  • Unknown
Political stability of top producer Unknown
Political stability of top reserve holder Unknown


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)
Unknown 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 Radium Podcast
Transcript :

Chemistry in its element: radium


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 the self illuminating story of element number 88. Here's Brian Clegg.

Brian Clegg

There's something about Radium that is deliciously Victorian. It's not just that this radioactive element was discovered at the end of the Victorian era in 1898. There's also something about its early use as a universal restorative that has a peculiarly period feel. It was seen as a source of energy and brightness, it was included in toothpastes and quack potions - it was even rubbed into the scalp as a hair restorer.

But the application of radium that would bring it notoriety was its use in glow-in-the-dark paint. Frequently used to provide luminous readouts on clocks and watches, aircraft switches and instrument dials, the eerie blue glow of radium was seen as a harmless, practical source of night time illumination. It was only when a number of the workers who painted the luminous dials began to suffer from sores, anaemia and cancers around the mouth that it was realized that something was horribly wrong. The women workers would regularly bring their paintbrushes to a point by licking them. This left enough radioactive residue in their mouths to cause cell damage. Eventually over 100 of the workers would die from the effects.

A more famous victim of radium was its discoverer, the double Nobel prize winner Marie Curie, born Maria Sklodowska. Working with her husband Pierre, Marie Curie was studying pitchblende, a mineral from North Bohemia that contained uranium. Pitchblende was mined near what's now Jachymov in the Czech Republic, and after the uranium had been extracted to be used to colour pottery glazes and tint photographs, the residual slag was dumped in a nearby forest. Without the uranium, the pitchblende proved still to be radioactive - in fact whatever the other radioactive material was, it was much more radioactive than the uranium itself.

Marie Curie wrote to sister Bronia that 'The radiation that I couldn't explain comes from a new chemical element. The element is there and I've got to find it! We are sure!' After working through tonnes of the pitchblende slag, the Curies identified two new elements in the remaining material - polonium and radium. They finally isolated radium in 1902 in its pure metal form. Radium was named for the Latin for a ray and proved to be the most radioactive natural substance ever discovered.

Although Marie Curie lived until 1934, her death from aplastic anaemia is almost certainly due to her exposure to radioactive materials, particularly radium. To this day her notebooks and papers have to be kept in lead lined boxes and handled with protective clothing, as they remain radioactive.

Radium occurs naturally as uranium decays - though only in very small quantities. It took many tonnes of pitchblende to produce the tenth of a gram of radium that the Curies eventually extracted. It's classified in the periodic table as an alkaline earth metal - the heaviest of the series - putting it alongside more familiar metals like magnesium and calcium. With atomic number 88, it has four natural isotopes of atomic weight 228, 226, 224 and 223 - though there are a remarkable 21 more artificial isotopes.

A later starring role for radium would be as the source of alpha particles - helium nuclei - used by Rutherford in 1909 at the Cavendish laboratory in Cambridge to fire at a thin gold foil. Radium decays to radon, throwing out an alpha particle from its nucleus. Unexpectedly, Rutherford's assistants Hans Geiger and Ernest Marsden found that a very few of the alpha particles bounced back - Rutherford likened it to 'firing a 15 inch shell at a piece of tissue paper and having it come back and hit you.' This behaviour was used to deduce the existence of a compact, dense nucleus in the atom - radium proved the key to unlocking the atom's structure.

Radium's main practical use has been in medicine, producing radon gas from radium chloride to be used in radiotherapy for cancer. This was a process started in Marie Curie's time. The early researchers found they received skin burns from handling the radioactive materials, and when the Curies worked with doctors, they discovered that radiation could be used to reduce or even cure tumours. This became known as Curie therapy, and the Sorbonne in Paris set up a laboratory partly for Curie to continue her research, and partly to study the medical applications of radiation, which would become known as the Radium Institute.

If you were to hold a piece of radium in your hand, it would feel warm. Initially a bright white, it would blacken as it reacted with the air to form radium nitride. It would stay solid - radium doesn't melt until around 700 degrees Celsius. It would also crackle and spit on the surface of your palm as it reacted with the water on your skin to produce radium hydroxide. Holding radium not something I'd recommend, though. Radium is constantly decaying, producing the alpha particles Rutherford used, beta particles, which are fast electrons, and gamma rays, like high energy X-rays, which would be slamming through your flesh, disrupting the DNA and causing cellular damage. The isotopes of radium vary in half life - the time it takes for half the molecules in a sample to delay - from 1,602 years for the most stable isotope, radium 226, to 11½ days for radium 223.

This is an element to be handled with care. Yet for anyone brought up on children's fiction full of ray guns and in a world were there were still X-ray machines to check your shoe size, it has a nostalgic feel that will ever make it fascinating.

Chris Smith

One wonders whether the podcasters of next century will be talking the same way about mobile phones, microwave ovens and MRI scanners. That was Bristol based science writer Brian Clegg with the story of radium. Next week to a metal capable of terrible cruelty to cancer.

Katherine Haxton

In the early 1960s, Barnett Rosenberg was conducting experiments on bacteria, measuring the effects of electrical currents on cell growth. The E.coli bacteria were abnormally long during the experiment, something that could not be attributed to the electric current. A number of platinum compounds were being formed due to reaction of the buffer and the platinum electrode. Cisplatin was found to inhibit cell division thus causing the elongation of the bacteria and was tested in was tested in mice for anticancer properties. Cisplatin today is widely used to treat epithelial malignancies with outstanding results in the treatment of testicular cancers.

Chris Smith

So we've got overgrown E.coli to blame for the discovery of platinum based anti cancer compounds. And you can find out how all of that came about with Keele University's Katherine Haxton on next week's Chemistry in its element. I'm Chris Smith, thank you for listening and for this week 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 1998-2017.



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 4.1), 2015, National Institute of Standards and Technology, Gaithersburg, MD, accessed November 2016.
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

Elements 1-112, 114, 116 and 117 © John Emsley 2012. Elements 113, 115, 117 and 118 © Royal Society of Chemistry 2017.



Produced by The Naked Scientists.


Periodic Table of Videos

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