Radium: Properties, Applications, and Safety
Ra Element
Radium is a type of chemical element that has the symbol Ra and an atomic number of 88. It is the sixth element in the group 2 of the periodic table, also called the alkaline earth metals. Pure radium is silvery-white in colour, however, it readily reacts with nitrogen rather than oxygen when exposed to air, and forms a black surface layer of radium nitride. All of the isotopes of radium are highly radioactive and the most stable isotope is radium-226 that has a half-life of 1600 years and decays into radon gas (specifically the isotope called radon-222). When the element radium decays, it yields ionizing radiation as a product that can excite the fluorescent chemicals and cause radioluminescence. In this article, we will learn about radium, the Ra element in detail, the use of radium, radium properties, the radium electronic configuration, and the effects of radium.
What is Radium?
Radium is a type of chemical element with a symbol Ra. It is the sixth element that lies in the group 2 of the periodic table. Pure form of radium is silvery-white in colour, however, it combines with nitrogen readily when it is exposed to air and forms a black surface layer of the radium nitride. Radium was discovered in the year 1898 by Marie Sklodowska Curie and Perre Curie in the form of radium chloride. They had extracted the radium compound from the element called uraninite. It is found in the uranium ores at the concentration of 1 part per 3 million parts uranium.
Physical Properties of Radium
Let us now look at the physical properties of radium.
Radium is known to be the heaviest known alkaline earth metal and is the one and only radioactive member of its periodic group. Its physical and chemical properties are much closely similar to its lighter congener which is barium.
Radium is a highly reactive metal to be known and it always exhibits its group oxidation state. It has the tendency to form the colourless Ra²⁺ cation in aqueous solution, which is highly basic in nature and does not form any complexes. Most of the radium compounds are for this reason simple ionic compounds.
Radium emits the alpha rays, beta rays, and gamma rays when it is mixed with the beryllium produces neutrons.
Let us now take a look at some of the chemical properties of radium.
Chemical Properties of Radium
Radium Uses
Let us now take a look at what can radium be used for and see some of its applications.
Some of the practical uses of radium are due to its radioactive properties. Radium was previously used in the self-luminous paints for watches, aircraft switches, nuclear panels clocks, and instrument dials.
Ra was earlier used as an additive in the products like hair cream, toothpaste, and even food items.
Radium was also used in the field of medicine for producing radon gas that in turn was used as a cancer treatment.
Health Effects of Radium
Now that you know about radium, let us see how radium is harmful to the health of the humans.
Radium is highly radiotoxic and carcinogen when it is inhaled, ingested or exposed and when it is used in the treatment of cancer and several other body disorders. The Ra element is more than a million times more radioactive than the same mass as that of uranium.
FAQs on Radium: Properties, Applications, and Safety
1. What are the fundamental properties of radium?
Radium (Ra) is a silvery-white, radioactive alkaline earth metal. Its fundamental properties include:
- Atomic Structure: An atomic number of 88, meaning it has 88 protons and 88 electrons. Its most stable isotope, Radium-226, has 138 neutrons.
- Physical State: It is a solid at room temperature with a density of 5.5 g/cm³.
- Reactivity: It is the heaviest and most reactive of the alkaline earth metals, tarnishing quickly when exposed to air.
- Chemical Behaviour: It consistently shows a +2 oxidation state in its compounds, similar to other Group 2 elements like barium.
2. Where is radium naturally found and why is it so rare?
Radium is found in trace amounts in uranium and thorium ores, such as uraninite. It is not found as a free element due to its high reactivity. Radium is exceptionally rare because all its isotopes are radioactive and have relatively short half-lives. For instance, its most stable isotope, Radium-226, has a half-life of 1600 years. As it forms from the decay of uranium, it also decays into other elements, preventing it from accumulating in large quantities.
3. What were the historical applications of radium, and why are they no longer used?
Historically, radium was used in various applications due to its radioluminescence (glowing property):
- Self-luminous paints: Used for watch dials, aircraft switches, and clock faces to make them visible in the dark.
- Medicine: It was once used in radiotherapy for cancer and was even added to consumer products like toothpaste and water under the false belief that it had curative powers.
4. What are the primary safety concerns and health effects associated with radium exposure?
Radium is extremely hazardous due to its intense radioactivity. The primary safety concerns are:
- Internal Exposure: If ingested or inhaled, the body treats radium like calcium, incorporating it into bones. The close-range alpha particle emission damages bone marrow and surrounding tissues, leading to a high risk of bone cancer and leukaemia.
- External Exposure: Gamma rays emitted by radium can travel long distances and penetrate the body, increasing cancer risk.
- Radon Gas: Radium decays into radon, a radioactive gas. Inhaling radon is a leading cause of lung cancer.
5. How does the reactivity of radium compare to other alkaline earth metals like magnesium and barium?
Radium is the most reactive alkaline earth metal, following the periodic trend for Group 2 where reactivity increases down the group. This is because the atomic radius increases and ionisation energy decreases, making it easier to lose valence electrons.
- Magnesium (Mg): Reacts slowly with cold water.
- Barium (Ba): Reacts vigorously with water.
- Radium (Ra): Reacts even more violently with water than barium to form radium hydroxide and hydrogen gas. Its high reactivity is due to its large atomic size and the ease with which it loses its two valence electrons.
6. What is the difference between the glow of radium paint and phosphorescence?
The glow from radium paint and phosphorescence are fundamentally different processes:
- Radium Paint (Radioluminescence): The glow is continuous and self-powered. Radium's radioactive decay emits particles that strike a fluorescent material (like zinc sulphide) in the paint, causing it to emit light. This process does not require any external light source to "charge" it.
- Phosphorescence: This is the process seen in common "glow-in-the-dark" items. A phosphorescent material absorbs energy from an external light source (like a lamp) and then slowly re-emits it as light. The glow fades as the stored energy is depleted and it needs to be recharged with light.
7. How does radium's electron configuration explain its chemical behaviour?
Radium's electron configuration is [Rn] 7s². This configuration is key to understanding its chemical properties:
- Valence Electrons: It has two electrons in its outermost shell (the 7s orbital). It readily loses these two electrons to achieve the stable electron configuration of the noble gas Radon (Rn).
- Oxidation State: By losing two electrons, it forms a cation with a +2 charge (Ra²⁺). This explains why radium almost exclusively exhibits a +2 oxidation state in its compounds.
- Group 2 Behaviour: This tendency to lose two electrons is characteristic of all alkaline earth metals, which is why radium shares chemical similarities with barium and calcium.
8. Why is Radium-226 significant in the study of radioactivity?
Radium-226 is historically and scientifically significant for several reasons. It was through studying radium that Marie and Pierre Curie developed fundamental concepts of radioactivity. Its relatively long half-life (1600 years) and its decay into the radioactive gas radon (Rn-222) made it a crucial element for early research into nuclear physics, decay chains, and the biological effects of radiation. It also served as the original standard for the unit of radioactivity, the curie (Ci).
