Radium is a chemical element with symbol Ra and atomic
number 88. It is the sixth element in group 2 of the periodic
table, also known as the alkaline earth metals. Pure radium is
silvery-white, but it readily reacts with nitrogen (rather than
oxygen) on exposure to air, forming a black surface layer of radium
nitride (Ra3N2). All isotopes of radium are highly radioactive, with
the most stable isotope being radium-226, which has a half-life of
1600 years and decays into radon gas (specifically the isotope
radon-222). When radium decays, ionizing radiation is a product, which
can excite fluorescent chemicals and cause radioluminescence.
Radium, in the form of radium chloride, was discovered by Marie and
Pierre Curie in 1898. They extracted the radium compound from
uraninite and published the discovery at the French Academy of
Sciences five days later.
Radium was isolated in its metallic state by
Marie Curie and
André-Louis Debierne through the electrolysis of
radium chloride in 1911.
In nature, radium is found in uranium and (to a lesser extent) thorium
ores in trace amounts as small as a seventh of a gram per ton of
Radium is not necessary for living organisms, and adverse
health effects are likely when it is incorporated into biochemical
processes because of its radioactivity and chemical reactivity.
Currently, other than its use in nuclear medicine, radium has no
commercial applications; formerly, it was used as a radioactive source
for radioluminescent devices and also in radioactive quackery for its
supposed curative powers. Today, these former applications are no
longer in vogue because radium's toxicity has since become known, and
less dangerous isotopes are used instead in radioluminescent devices.
1 Bulk properties
5.1 Historical applications
5.1.1 Luminescent paint
5.1.2 Commercial use
5.1.3 Medical use
7 Modern applications
9 See also
13 Further reading
14 External links
Radium is the heaviest known alkaline earth metal and is the only
radioactive member of its group. Its physical and chemical properties
most closely resemble its lighter congener barium.
Pure radium is a volatile silvery-white metal, although its lighter
congeners calcium, strontium, and barium have a slight yellow tint.
Its color rapidly vanishes in air, yielding a black layer of radium
nitride (Ra3N2). Its melting point is either 700 °C
(1,292 °F) or 960 °C (1,760 °F)[a] and its boiling
point is 1,737 °C (3,159 °F). Both of these values are
slightly lower than those of barium, confirming periodic trends down
the group 2 elements. Like barium and the alkali metals, radium
crystallizes in the body-centered cubic structure at standard
temperature and pressure: the radium–radium bond distance is
Radium has a density of 5.5 g/cm3,
higher than that of barium, again confirming periodic trends; the
radium-barium density ratio is comparable to the radium-barium atomic
mass ratio, due to the two elements' similar crystal
Main article: Isotopes of radium
Decay chain of 238U, the primordial progenitor of 226Ra
Radium has 33 known isotopes, with mass numbers from 202 to 234: all
of them are radioactive. Four of these – 223Ra (half-life
11.4 days), 224Ra (3.64 days), 226Ra (1600 years), and
228Ra (5.75 years) – occur naturally in the decay chains of
primordial thorium-232, uranium-235, and uranium-238 (223Ra from
uranium-235, 226Ra from uranium-238, and the other two from
thorium-232). These isotopes nevertheless still have half-lives too
short to be primordial radionuclides and only exist in nature from
these decay chains. Together with the artificial 225Ra (15 d),
these are the five most stable isotopes of radium. All other known
radium isotopes have half-lives under two hours, and the majority have
half-lives under a minute. At least 12 nuclear isomers have been
reported; the most stable of them is radium-205m, with a half-life of
between 130 and 230 milliseconds, which is still shorter than
twenty-four ground-state radium isotopes.
In the early history of the study of radioactivity, the different
natural isotopes of radium were given different names. In this scheme,
223Ra was named actinium X (AcX), 224Ra thorium X (ThX), 226Ra radium
(Ra), and 228Ra mesothorium 1 (MsTh1). When it was realized that
all of these are isotopes of the same element, many of these names
fell out of use, and "radium" came to refer to all isotopes, not just
226Ra. Some of radium-226's decay products received historical
names including "radium", ranging from radium A to radium G, with the
letter indicating approximately how far they were down the chain from
their parent 226Ra.
226Ra is the most stable isotope of radium and is the last isotope in
the (4n + 2) decay chain of uranium-238 with a half-life of
over a millennium: it makes up almost all of natural radium. Its
immediate decay product is the dense radioactive noble gas radon,
which is responsible for much of the danger of environmental
radium. It is 2.7 million times more radioactive than the same
molar amount of natural uranium (mostly uranium-238), due to its
proportionally shorter half-life.
A sample of radium metal maintains itself at a higher temperature than
its surroundings because of the radiation it emits – alpha
particles, beta particles, and gamma rays. More specifically, natural
radium (which is mostly 226Ra) emits mostly alpha particles, but other
steps in its decay chain (the uranium or radium series) emit alpha or
beta particles, and almost all particle emissions are accompanied by
Radium, like barium, is a highly reactive metal and always exhibits
its group oxidation state of +2. It forms the colorless Ra2+ cation
in aqueous solution, which is highly basic and does not form complexes
readily. Most radium compounds are therefore simple ionic
compounds, though participation from the 6s and 6p electrons (in
addition to the valence 7s electrons) is expected due to relativistic
effects and would enhance the covalent character of radium compounds
such as RaF2 and RaAt2. For this reason, the standard electrode
potential for the half-reaction Ra2+ (aq) + 2e− → Ra (s) is
−2.916 V, even slightly lower than the value −2.92 V for
barium, whereas the values had previously smoothly increased down the
group (Ca: −2.84 V; Sr: −2.89 V; Ba:
−2.92 V). The values for barium and radium are almost
exactly the same as those of the heavier alkali metals potassium,
rubidium, and caesium.
Solid radium compounds are white as radium ions provide no specific
coloring, but they gradually turn yellow and then dark over time due
to self-radiolysis from radium's alpha decay. Insoluble radium
compounds coprecipitate with all barium, most strontium, and most lead
Radium oxide (RaO) has not been characterized well past its existence,
despite oxides being common compounds for the other alkaline earth
Radium hydroxide (Ra(OH)2) is the most readily soluble among
the alkaline earth hydroxides and is a stronger base than its barium
congener, barium hydroxide. It is also more soluble than actinium
hydroxide and thorium hydroxide: these three adjacent hydroxides may
be separated by precipitating them with ammonia.
Radium chloride (RaCl2) is a colorless, luminous compound. It becomes
yellow after some time due to self-damage by the alpha radiation given
off by radium when it decays. Small amounts of barium impurities give
the compound a rose color. It is soluble in water, though less so
than barium chloride, and its solubility decreases with increasing
concentration of hydrochloric acid. Crystallization from aqueous
solution gives the dihydrate RaCl2·2H2O, isomorphous with its barium
Radium bromide (RaBr2) is also a colorless, luminous compound. In
water, it is more soluble than radium chloride. Like radium chloride,
crystallization from aqueous solution gives the dihydrate RaBr2·2H2O,
isomorphous with its barium analog. The ionizing radiation emitted by
radium bromide excites nitrogen molecules in the air, making it glow.
The alpha particles emitted by radium quickly gain two electrons to
become neutral helium, with builds up inside and weakens radium
bromide crystals. This effect sometimes causes the crystals to break
or even explode.
Radium nitrate (Ra(NO3)2) is a white compound that can be made by
dissolving radium carbonate in nitric acid. As the concentration of
nitric acid increases, the solubility of radium nitrate decreases, an
important property for the chemical purification of radium.
Radium forms much the same insoluble salts as its lighter congener
barium: it forms the insoluble sulfate (RaSO4, the most insoluble
known sulfate), chromate (RaCrO4), carbonate (RaCO3), iodate
(Ra(IO3)2), tetrafluoroberyllate (RaBeF4), and nitrate (Ra(NO3)2).
With the exception of the carbonate, all of these are less soluble in
water than the corresponding barium salts, but they are all
isostructural to their barium counterparts. Additionally, radium
phosphate, oxalate, and sulfite are probably also insoluble, as they
coprecipitate with the corresponding insoluble barium salts. The
great insolubility of radium sulfate (at 20 °C, only 2.1 mg
will dissolve in 1 kg of water) means that it is one of the less
biologically dangerous radium compounds. The large ionic radius of
Ra2+ (148 pm) results in weak complexation and poor extraction of
radium from aqueous solutions when not at high pH.
All isotopes of radium have half-lives much shorter than the age of
the Earth, so that any primordial radium would have decayed long ago.
Radium nevertheless still occurs in the environment, as the isotopes
223Ra, 224Ra, 226Ra, and 228Ra are part of the decay chains of natural
thorium and uranium isotopes; since thorium and uranium have very long
half-lives, these daughters are continually being regenerated by their
decay. Of these four isotopes, the longest-lived is 226Ra
(half-life 1600 years), a decay product of natural uranium.
Because of its relative longevity, 226Ra is the most common isotope of
the element, making up about one part per trillion of the Earth's
crust; essentially all natural radium is 226Ra. Thus, radium is
found in tiny quantities in the uranium ore uraninite and various
other uranium minerals, and in even tinier quantities in thorium
minerals. One ton of pitchblende typically yields about one seventh of
a gram of radium. One kilogram of the
Earth's crust contains about
900 picograms of radium, and one liter of sea water contains
about 89 femtograms of radium.
Pierre Curie experimenting with radium, a drawing by André
Glass tube of radium chloride kept by the US Bureau of Standards that
served as the primary standard of radioactivity for the United States
Marie Curie § New elements
Radium was discovered by Marie Sklodowska-
Curie and her husband Pierre
Curie on 21 December 1898, in a uraninite (pitchblende) sample.
While studying the mineral earlier, the Curies removed uranium from it
and found that the remaining material was still radioactive. They
separated out an element similar to bismuth from pitchblende in July
1898, that turned out to be polonium. They then separated out a
radioactive mixture consisting mostly of two components: compounds of
barium, which gave a brilliant green flame color, and unknown
radioactive compounds which gave carmine spectral lines that had never
been documented before. The Curies found the radioactive compounds to
be very similar to the barium compounds, except that they were more
insoluble. This made it possible for the Curies to separate out the
radioactive compounds and discover a new element in them. The Curies
announced their discovery to the
French Academy of Sciences
French Academy of Sciences on 26
December 1898. The naming of radium dates to about 1899, from
the French word radium, formed in Modern Latin from radius (ray): this
was in recognition of radium's power of emitting energy in the form of
In 1910, radium was isolated as a pure metal by
Marie Curie and
André-Louis Debierne through the electrolysis of a pure radium
chloride (RaCl2) solution using a mercury cathode, producing a
radium–mercury amalgam. This amalgam was then heated in an
atmosphere of hydrogen gas to remove the mercury, leaving pure radium
metal. The same year, E. Eoler isolated radium by thermal
decomposition of its azide, Ra(N3)2.
Radium metal was first
industrially produced in the beginning of the 20th century by Biraco,
a subsidiary company of
Union Minière du Haut Katanga
Union Minière du Haut Katanga (UMHK) in its
Olen plant in Belgium.
The common historical unit for radioactivity, the curie, is based on
the radioactivity of 226Ra.
Self-luminous white paint which contains radium on the face and hand
of an old clock.
Radium watch hands under ultraviolet light
Radium was formerly used in self-luminous paints for watches, nuclear
panels, aircraft switches, clocks, and instrument dials. A typical
self-luminous watch that uses radium paint contains around 1 microgram
of radium. In the mid-1920s, a lawsuit was filed against the
United States Radium Corporation
United States Radium Corporation by five dying "
Radium Girls" dial
painters who had painted radium-based luminous paint on the dials of
watches and clocks. The dial painters routinely licked their brushes
to give them a fine point, thereby ingesting radium. Their
exposure to radium caused serious health effects which included sores,
anemia, and bone cancer. This is because radium is treated as calcium
by the body, and deposited in the bones, where radioactivity degrades
marrow and can mutate bone cells.
During the litigation, it was determined that the company's scientists
and management had taken considerable precautions to protect
themselves from the effects of radiation, yet had not seen fit to
protect their employees. Additionally, for several years the companies
had attempted to cover up the effects and avoid liability by insisting
Radium Girls were instead suffering from syphilis. This
complete disregard for employee welfare had a significant impact on
the formulation of occupational disease labor law.
As a result of the lawsuit, the adverse effects of radioactivity
became widely known, and radium-dial painters were instructed in
proper safety precautions and provided with protective gear. In
particular, dial painters no longer licked paint brushes to shape them
(which caused some ingestion of radium salts).
Radium was still used
in dials as late as the 1960s, but there were no further injuries to
dial painters. This highlighted that the harm to the
could easily have been avoided.
From the 1960s the use of radium paint was discontinued. In many cases
luminous dials were implemented with non-radioactive fluorescent
materials excited by light; such devices glow in the dark after
exposure to light, but the glow fades. Where long-lasting
self-luminosity in darkness was required, safer radioactive
promethium-147 (half-life 2.6 years) or tritium (half-life 12 years)
paint was used; both continue to be used today. These had the
added advantage of not degrading the phosphor over time, unlike
Tritium emits very low-energy beta radiation (even
lower-energy than the beta radiation emitted by promethium) which
cannot penetrate the skin, rather than the penetrating gamma
radiation of radium and is regarded as safer.
Clocks, watches, and instruments dating from the first half of the
20th century, often in military applications, may have been painted
with radioactive luminous paint. They are usually no longer luminous;
however, this is not due to radioactive decay of the radium (which has
a half-life of 1600 years) but to the fluorescence of the zinc sulfide
fluorescent medium being worn out by the radiation from the
radium. The appearance of an often thick layer of green or
yellowish brown paint in devices from this period suggests a
radioactive hazard. The radiation dose from an intact device is
relatively low and usually not an acute risk; but the paint is
dangerous if released and inhaled or ingested.
Hotel postcard advertising radium baths, c.1940s
Radium was once an additive in products such as toothpaste, hair
creams, and even food items due to its supposed curative powers.
Such products soon fell out of vogue and were prohibited by
authorities in many countries after it was discovered they could have
serious adverse health effects. (See, for instance,
Revigator types of "
Radium water" or "Standard
Radium Solution for
Drinking".) Spas featuring radium-rich water are still
occasionally touted as beneficial, such as those in Misasa, Tottori,
Japan. In the U.S., nasal radium irradiation was also administered to
children to prevent middle-ear problems or enlarged tonsils from the
late 1940s through the early 1970s.
Radium (usually in the form of radium chloride or radium bromide) was
used in medicine to produce radon gas which in turn was used as a
cancer treatment; for example, several of these radon sources were
used in Canada in the 1920s and 1930s. However, many
treatments that were used in the early 1900s are not used anymore
because of the harmful effects radium bromide exposure caused. Some
examples of these effects are anaemia, cancer, and genetic
mutations. Safer gamma emitters such as 60Co, which is less costly
and available in larger quantities, are usually used today to replace
the historical use of radium in this application.
Early in the 1900s, biologists used radium to induce mutations and
study genetics. As early as 1904, Daniel MacDougal used radium in an
attempt to determine whether it could provoke sudden large mutations
and cause major evolutionary shifts.
Thomas Hunt Morgan
Thomas Hunt Morgan used radium to
induce changes resulting in white-eyed fruit flies. Nobel-winning
biologist Hermann Muller briefly studied the effects of radium on
fruit fly mutations before turning to more affordable x-ray
Howard Atwood Kelly, one of the founding physicians of Johns Hopkins
Hospital, was a major pioneer in the medical use of radium to treat
cancer. His first patient was his own aunt in 1904, who died
shortly after surgery. Kelly was known to use excessive amounts of
radium to treat various cancers and tumors. As a result, some of his
patients died from radium exposure. His method of radium
application was inserting a radium capsule near the affected area,
then sewing the radium "points" directly to the tumor. This was
the same method used to treat Henrietta Lacks, the host of the
original HeLa cells, for cervical cancer. Currently, safer and
more available radioisotopes are used instead.
Uranium had no large scale application in the late 19th century and
therefore no large uranium mines existed. In the beginning the only
large source for uranium ore was the silver mines in Joachimsthal,
Austria-Hungary (now Jáchymov, Czech Republic). The uranium ore
was only a byproduct of the mining activities.
In the first extraction of radium
Curie used the residues after
extraction of uranium from pitchblende. The uranium had been extracted
by dissolution in sulfuric acid leaving radium sulfate, which is
similar to barium sulfate but even less soluble in the residues. The
residues also contained rather substantial amounts of barium sulfate
which thus acted as a carrier for the radium sulfate. The first steps
of the radium extraction process involved boiling with sodium
hydroxide followed by hydrochloric acid treatment to remove as much as
possible of other compounds. The remaining residue was then treated
with sodium carbonate to convert the barium sulfate into barium
carbonate carrying the radium, thus making it soluble in hydrochloric
acid. After dissolution the barium and radium are reprecipitated as
sulfates and this was repeated one or few times, for further
purification of the mixed sulfate. Some impurities, that form
insoluble sulfides, were removed by treating the chloride solution
with hydrogen sulfide followed by filtering. When the mixed sulfate
were pure enough they were once more converted to mixed chloride and
barium and radium were separated by fractional crystallisation while
monitoring the progress using a spectroscope (radium gives
characteristic red lines in contrast to the green barium lines), and
After the isolation of radium by Marie and
Pierre Curie from uranium
ore from Joachimsthal several scientists started to isolate radium in
small quantities. Later small companies purchased mine tailings from
Joachimsthal mines and started isolating radium. In 1904 the Austrian
government nationalised the mines and stopped exporting raw ore. For
some time the radium availability was low.
The formation of an Austrian monopoly and the strong urge of other
countries to have access to radium led to a worldwide search for
uranium ores. The United States took over as leading producer in the
early 1910s. The
Carnotite sands in
Colorado provide some of the
element, but richer ores are found in the Congo and the area of the
Great Bear Lake
Great Bear Lake and the
Great Slave Lake
Great Slave Lake of northwestern
Canada. Neither of the deposits is mined for radium but the
uranium content makes mining profitable.
The Curies' process was still used for industrial radium extraction in
1940, but mixed bromides were then used for the fractionation. If
the barium content of the uranium ore is not high enough it is easy to
add some to carry the radium. These processes were applied to high
grade uranium ores but may not work well with low grade ores.
Small amounts of radium were still extracted from uranium ore by this
method of mixed precipitation and ion exchange as late as the
1990s, but today they are extracted only from spent nuclear
fuel. In 1954, the total worldwide supply of purified radium
amounted to about 5 pounds (2.3 kg) and it is still in this
range today, while the annual production of pure radium compounds is
only about 100 g in total today. The chief radium-producing
countries are Belgium, Canada, the Czech Republic, Slovakia, the
United Kingdom, and Russia. The amounts of radium produced were
and are always relatively small; for example, in 1918, 13.6 g of
radium were produced in the United States.  The metal is isolated
by reducing radium oxide with aluminium metal in a vacuum at
Some of the few practical uses of radium are derived from its
radioactive properties. More recently discovered radioisotopes, such
as cobalt-60 and caesium-137, are replacing radium in even these
limited uses because several of these isotopes are more powerful
emitters, safer to handle, and available in more concentrated
The isotope 223Ra (under the trade name Xofigo) was approved by the
Food and Drug Administration
Food and Drug Administration in 2013 for use in medicine
as a cancer treatment of bone metastasis. The main indication
of treatment with
Xofigo is the therapy of bony metastases from
castration-resistant prostate cancer due to the favourable
characteristics of this alpha-emitter radiopharmaceutical. 225Ra
has also been used in experiments concerning therapeutic irradiation,
as it is the only reasonably long-lived radium isotope which does not
have radon as one of its daughters.
Radium is still used today as a radiation source in some industrial
radiography devices to check for flawed metallic parts, similarly to
X-ray imaging. When mixed with beryllium, radium acts as a neutron
source. Radium-beryllium neutron sources are still sometimes
used even today, but other materials such as polonium are now
more common: about 1500 polonium-beryllium neutron sources, with an
individual activity of 1,850 Ci (68 TBq), have been used
annually in Russia. These RaBeF4-based (α, n) neutron sources
have been deprecated despite the high number of neutrons they emit
(1.84×106 neutrons per second) in favour of 241Am–Be sources.
Today, the isotope 226Ra is mainly used to form 227Ac by neutron
irradiation in a nuclear reactor.
Radium is highly radioactive and its immediate daughter, radon gas, is
also radioactive. When ingested, 80% of the ingested radium leaves the
body through the feces, while the other 20% goes into the bloodstream,
mostly accumulating in the bones. Exposure to radium, internal or
external, can cause cancer and other disorders, because radium and
radon emit alpha and gamma rays upon their decay, which kill and
mutate cells. At the time of the
Manhattan Project in 1944, the
"tolerance dose" for workers was set at 0.1 micrograms of ingested
Some of the biological effects of radium were apparent from the start.
The first case of so-called "radium-dermatitis" was reported in 1900,
only 2 years after the element's discovery. The French physicist
Antoine Becquerel carried a small ampoule of radium in his waistcoat
pocket for 6 hours and reported that his skin became ulcerated. Pierre
Marie Curie were so intrigued by radiation that they sacrificed
their own health to learn more about it.
Pierre Curie attached a tube
filled with radium to his arm for ten hours, which resulted in the
appearance of a skin lesion, suggesting the use of radium to attack
cancerous tissue as it had attacked healthy tissue. Handling of
radium has been blamed for Marie Curie's death due to aplastic anemia.
A significant amount of radium's danger comes from its daughter radon:
being a gas, it can enter the body far more readily than can its
Today, 226Ra is considered to be the most toxic of the quantity
radioelements, and it must be handled in tight glove boxes with
significant airstream circulation that is then treated to avoid escape
of its daughter 222Rn to the environment. Old ampoules containing
radium solutions must be opened with care because radiolytic
decomposition of water can produce an overpressure of hydrogen and
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Find more aboutRadiumat's sister projects
Definitions from Wiktionary
Media from Wikimedia Commons
Learning resources from Wikiversity
"Lateral Science: The Discovery of Radium". July 8, 2012. Archived
from the original on March 9, 2016. Retrieved 13 May 2017.
Radium Water Bath in Oklahoma
NLM Hazardous Substances Databank – Radium, Radioactive
Annotated bibliography for radium from the Alsos Digital Library for
The Poisoner Next Door – Japan Today, 10/20/2001
The Periodic Table of Videos
The Periodic Table of Videos (University of Nottingham)
Radioactivity.eu.com (Created and maintained by physicists)
Periodic table (Large cells)
Alkaline earth metal