| Discovery date | 1898 |
| Discovered by | Pierre and Marie Curie |
| Origin of the name | The name is derived from the Latin 'radius', meaning ray. |
| Allotropes |
|
Ra
Radium
88
[226]
| Group | 2 | Melting point | 696°C, 1285°F, 969 K |
| Period | 7 | Boiling point | 1500°C, 2732°F, 1773 K |
| Block | s | Density (g cm−3) | 5 |
| 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
The image represents the former use of radium in luminous paint used for clock and watch dials.
Appearance
A soft, shiny and silvery radioactive metal.
Uses
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.
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 (Å) | 2.83 | Covalent radius (Å) | 2.11 |
| Electron affinity (kJ mol−1) | 9.65 |
Electronegativity (Pauling scale) |
0.9 |
|
Ionisation energies (kJ mol−1) |
1st
509.29
2nd
979.051
3rd
-
4th
-
5th
-
6th
-
7th
-
8th
-
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| 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 y | sf | ||||
| 228Ra | 228.031 | - | 5.76 y | β- | |
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βf | ||||
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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) |
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| Pressure (Pa) |
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| Listen to Radium Podcast |
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Transcript :
Chemistry in its element: radium (Promo) You're listening to Chemistry in its element brought to you by Chemistry World, the magazine of the Royal Society of Chemistry. (End promo) 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. (Promo) Chemistry in its element is brought to you by the Royal Society of Chemistry and produced by thenakedscientists.com. There's more information and other episodes of Chemistry in its element on our website at chemistryworld.org/elements. (End promo)
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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.
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.
Derived in part from material provided by the British Geological Survey © NERC.
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.
Created by video journalist Brady Haran working with chemists at The University of Nottingham.
© Murray Robertson 1998-2017.
Data
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.
Podcasts
Produced by The Naked Scientists.
Periodic Table of Videos
Created by video journalist Brady Haran working with chemists at The University of Nottingham.
