Radon is a chemical element; it has symbol Rn and atomic number 86. It is a radioactive noble gas and is colorless and odorless. Of the three naturally occurring radon isotopes, only radon-222 has a sufficiently long half-life (3.825 days) for it to be released from the soil and rock where it is generated. Radon isotopes are the immediate decay products of radium isotopes. The instability of radon-222, its most stable isotope, makes radon one of the rarest elements. Radon will be present on Earth for several billion more years despite its short half-life, because it is constantly being produced as a step in the decay chains of uranium-238 and thorium-232, of which both are abundant radioactive nuclides with half-lives of at least several billion years. The decay of radon produces many other short-lived nuclides, known as "radon daughters", ending at stable isotopes of lead.[1] Radon-222 occurs in significant quantities as a step in the normal radioactive decay chain of uranium-238, also known as the uranium series, which slowly decays into a variety of radioactive nuclides and eventually decays into stable lead-206. Radon-220 occurs in minute quantities as an intermediate step in the decay chain of thorium-232, also known as the thorium series, which eventually decays into stable lead-208.
Under standard conditions, radon is gaseous and can be easily inhaled, posing a health hazard. However, the primary danger comes not from radon itself, but from its decay products, known as radon daughters. These decay products, often existing as single atoms or ions, can attach themselves to airborne dust particles. Although radon is a noble gas and does not adhere to lung tissue (meaning it is often exhaled before decaying), the radon daughters attached to dust are more likely to stick to the lungs. This increases the risk of harm, as the radon daughters can cause damage to lung tissue.[2] Radon and its daughters are, taken together, often the single largest contributor to an individual's background radiation dose, but due to local differences in geology,[3] the level of exposure to radon gas differs by location. A common source of environmental radon is uranium-containing minerals in the ground; it therefore accumulates in subterranean areas such as basements. Radon can also occur in ground water, such as spring waters and hot springs.[4] Radon trapped in permafrost may be released by climate-change-induced thawing of permafrosts.[5] It is possible to test for radon in buildings, and to use techniques such as sub-slab depressurization for mitigation.[6] [7]
Epidemiological studies have shown a clear association between breathing high concentrations of radon and incidence of lung cancer.[8] [9] [10] Radon is a contaminant that affects indoor air quality worldwide. According to the United States Environmental Protection Agency (EPA), radon is the second most frequent cause of lung cancer, after cigarette smoking, causing 21,000 lung cancer deaths per year in the United States. About 2,900 of these deaths occur among people who have never smoked. While radon is the second most frequent cause of lung cancer, it is the number one cause among non-smokers, according to EPA policy-oriented estimates. Significant uncertainties exist for the health effects of low-dose exposures.[11]
Radon is a colorless, odorless, and tasteless[12] gas and therefore is not detectable by human senses alone. At standard temperature and pressure, it forms a monatomic gas with a density of 9.73 kg/m3, about 8 times the density of the Earth's atmosphere at sea level, 1.217 kg/m3.[13] It is one of the densest gases at room temperature (a few are denser, e.g. CF3(CF2)2CF3 and WF6) and is the densest of the noble gases. Although colorless at standard temperature and pressure, when cooled below its freezing point of, it emits a brilliant radioluminescence that turns from yellow to orange-red as the temperature lowers.[14] Upon condensation, it glows because of the intense radiation it produces.[15] It is sparingly soluble in water, but more soluble than lighter noble gases. It is appreciably more soluble in organic liquids than in water. Its solubility equation is as follows,[16] [17] [18]
\chi=\exp(B/T-A),
where
\chi
T
A
B
Radon is a member of the zero-valence elements that are called noble gases, and is chemically not very reactive. The 3.8-day half-life of radon-222 makes it useful in physical sciences as a natural tracer. Because radon is a gas at standard conditions, unlike its decay-chain parents, it can readily be extracted from them for research.
It is inert to most common chemical reactions, such as combustion, because the outer valence shell contains eight electrons. This produces a stable, minimum energy configuration in which the outer electrons are tightly bound.[19] Its first ionization energy—the minimum energy required to extract one electron from it—is 1037 kJ/mol.[20] In accordance with periodic trends, radon has a lower electronegativity than the element one period before it, xenon, and is therefore more reactive. Early studies concluded that the stability of radon hydrate should be of the same order as that of the hydrates of chlorine or sulfur dioxide, and significantly higher than the stability of the hydrate of hydrogen sulfide .[21]
Because of its cost and radioactivity, experimental chemical research is seldom performed with radon, and as a result there are very few reported compounds of radon, all either fluorides or oxides. Radon can be oxidized by powerful oxidizing agents such as fluorine, thus forming radon difluoride .[22] [23] It decomposes back to its elements at a temperature of above 523K, and is reduced by water to radon gas and hydrogen fluoride: it may also be reduced back to its elements by hydrogen gas. It has a low volatility and was thought to be . Because of the short half-life of radon and the radioactivity of its compounds, it has not been possible to study the compound in any detail. Theoretical studies on this molecule predict that it should have a Rn–F bond distance of 2.08 ångströms (Å), and that the compound is thermodynamically more stable and less volatile than its lighter counterpart xenon difluoride .[24] The octahedral molecule was predicted to have an even lower enthalpy of formation than the difluoride.[25] The [RnF]+ ion is believed to form by the following reaction:[26]
Rn (g) + 2 (s) → (s) + 2 (g)
For this reason, antimony pentafluoride together with chlorine trifluoride and have been considered for radon gas removal in uranium mines due to the formation of radon–fluorine compounds. Radon compounds can be formed by the decay of radium in radium halides, a reaction that has been used to reduce the amount of radon that escapes from targets during irradiation. Additionally, salts of the [RnF]+ cation with the anions,, and are known.[27] Radon is also oxidised by dioxygen difluoride to at 173K.
Radon oxides are among the few other reported compounds of radon;[28] only the trioxide has been confirmed. The higher fluorides and have been claimed and are calculated to be stable,[29] but their identification is unclear.[30] They may have been observed in experiments where unknown radon-containing products distilled together with xenon hexafluoride: these may have been,, or both. Trace-scale heating of radon with xenon, fluorine, bromine pentafluoride, and either sodium fluoride or nickel fluoride was claimed to produce a higher fluoride as well which hydrolysed to form . While it has been suggested that these claims were really due to radon precipitating out as the solid complex [RnF][NiF<sub>6</sub>]2−, the fact that radon coprecipitates from aqueous solution with has been taken as confirmation that was formed, which has been supported by further studies of the hydrolysed solution. That [RnO<sub>3</sub>F]− did not form in other experiments may have been due to the high concentration of fluoride used. Electromigration studies also suggest the presence of cationic [HRnO<sub>3</sub>]+ and anionic [HRnO<sub>4</sub>]− forms of radon in weakly acidic aqueous solution (pH > 5), the procedure having previously been validated by examination of the homologous xenon trioxide.
The decay technique has also been used. Avrorin et al. reported in 1982 that 212Fr compounds cocrystallised with their caesium analogues appeared to retain chemically bound radon after electron capture; analogies with xenon suggested the formation of RnO3, but this could not be confirmed.[31]
It is likely that the difficulty in identifying higher fluorides of radon stems from radon being kinetically hindered from being oxidised beyond the divalent state because of the strong ionicity of radon difluoride and the high positive charge on radon in RnF+; spatial separation of molecules may be necessary to clearly identify higher fluorides of radon, of which is expected to be more stable than due to spin–orbit splitting of the 6p shell of radon (RnIV would have a closed-shell 6s6p configuration). Therefore, while should have a similar stability to xenon tetrafluoride, would likely be much less stable than xenon hexafluoride : radon hexafluoride would also probably be a regular octahedral molecule, unlike the distorted octahedral structure of, because of the inert pair effect.[32] [33] Because radon is quite electropositive for a noble gas, it is possible that radon fluorides actually take on highly fluorine-bridged structures and are not volatile.[33] Extrapolation down the noble gas group would suggest also the possible existence of RnO, RnO2, and RnOF4, as well as the first chemically stable noble gas chlorides RnCl2 and RnCl4, but none of these have yet been found.
Radon carbonyl (RnCO) has been predicted to be stable and to have a linear molecular geometry.[34] The molecules and RnXe were found to be significantly stabilized by spin-orbit coupling.[35] Radon caged inside a fullerene has been proposed as a drug for tumors.[36] [37] Despite the existence of Xe(VIII), no Rn(VIII) compounds have been claimed to exist; should be highly unstable chemically (XeF8 is thermodynamically unstable). It is predicted that the most stable Rn(VIII) compound would be barium perradonate (Ba2RnO6), analogous to barium perxenate. The instability of Rn(VIII) is due to the relativistic stabilization of the 6s shell, also known as the inert pair effect.
Radon reacts with the liquid halogen fluorides ClF,,,,, and to form . In halogen fluoride solution, radon is nonvolatile and exists as the RnF+ and Rn2+ cations; addition of fluoride anions results in the formation of the complexes and, paralleling the chemistry of beryllium(II) and aluminium(III). The standard electrode potential of the Rn2+/Rn couple has been estimated as +2.0 V,[38] although there is no evidence for the formation of stable radon ions or compounds in aqueous solution.
See main article: Isotopes of radon.
Radon has no stable isotopes. Thirty-nine radioactive isotopes have been characterized, with mass numbers ranging from 193 to 231.[39] [40] Six of them, from 217 to 222 inclusive, occur naturally. The most stable isotope is Rn (half-life 3.82 days), which is a decay product of Ra, the latter being itself a decay product of U.[41] A trace amount of the (highly unstable) isotope Rn (half-life about 35 milliseconds) is also among the daughters of Rn. The isotope Rn would be produced by the double beta decay of natural Po; while energetically possible, this process has however never been seen.[42]
Three other radon isotopes have a half-life of over an hour: Rn (about 15 hours), Rn (2.4 hours) and Rn (about 1.8 hours). However, none of these three occur naturally. Rn, also called thoron, is a natural decay product of the most stable thorium isotope (Th). It has a half-life of 55.6 seconds and also emits alpha radiation. Similarly, Rn is derived from the most stable isotope of actinium (Ac)—named "actinon"—and is an alpha emitter with a half-life of 3.96 seconds. No radon isotopes occur significantly in the neptunium (Np) decay series, though trace amounts of the isotopes Rn (26 minutes) and Rn (0.5 millisecond) are produced in minor branches.
222Rn belongs to the radium and uranium-238 decay chain, and has a half-life of 3.8235 days. Its first four products (excluding marginal decay schemes) are very short-lived, meaning that the corresponding disintegrations are indicative of the initial radon distribution. Its decay goes through the following sequence:
The radon equilibrium factor[43] is the ratio between the activity of all short-period radon progenies (which are responsible for most of radon's biological effects), and the activity that would be at equilibrium with the radon parent.
If a closed volume is constantly supplied with radon, the concentration of short-lived isotopes will increase until an equilibrium is reached where the overall decay rate of the decay products equals that of the radon itself. The equilibrium factor is 1 when both activities are equal, meaning that the decay products have stayed close to the radon parent long enough for the equilibrium to be reached, within a couple of hours. Under these conditions, each additional pCi/L of radon will increase exposure by 0.01 working level (WL, a measure of radioactivity commonly used in mining). These conditions are not always met; in many homes, the equilibrium factor is typically 40%; that is, there will be 0.004 WL of daughters for each pCi/L of radon in the air. 210Pb takes much longer (decades) to come in equilibrium with radon, but, if the environment permits accumulation of dust over extended periods of time, 210Pb and its decay products may contribute to overall radiation levels as well.
Because of their electrostatic charge, radon progenies adhere to surfaces or dust particles, whereas gaseous radon does not. Attachment removes them from the air, usually causing the equilibrium factor in the atmosphere to be less than 1. The equilibrium factor is also lowered by air circulation or air filtration devices, and is increased by airborne dust particles, including cigarette smoke. The equilibrium factor found in epidemiological studies is 0.4.[44]