Perchlorate

A perchlorate is the name for a chemical compound containing the perchlorate ion, ClO
4
. The majority of perchlorates are produced commercially. Perchlorate salts are mainly used for propellants, exploiting properties as powerful oxidizing agents and to control static electricity in food packaging.[2] Perchlorate contamination in food, water and other parts of the environment has been studied in the U.S. because of its harmful effects on human health. Perchlorate reduces thyroid hormone production in the thyroid gland.

Most perchlorates are colorless solids that are soluble in water. Four perchlorates are of primary commercial interest: ammonium perchlorate (NH4ClO4), perchloric acid (HClO4), potassium perchlorate (KClO4), and sodium perchlorate (NaClO4). Perchlorate is the anion resulting from the dissociation of perchloric acid and its salts upon their dissolution in water. Many perchlorate salts are soluble in non-aqueous solutions.[3]

Perchlorate
Skeletal model of perchlorate showing various dimensions
Ball-and-stick model of the perchlorate ion
Spacefill model of perchlorate
Names
Systematic IUPAC name
Perchlorate[1]
Identifiers
3D model (JSmol)
ChEBI
ChEMBL
ChemSpider
DrugBank
ECHA InfoCard 100.152.366
2136
MeSH 180053
Properties
ClO
4
Molar mass 99.451 g mol−1
Conjugate acid Perchloric acid
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).

Production

Perchlorate salts are produced industrially by the oxidation of solutions of sodium chlorate by electrolysis. This method is used to prepare sodium perchlorate. The main application is for rocket fuel.[4] The reaction of perchloric acid with bases, such as ammonium hydroxide, give salts. The highly valued ammonium perchlorate can be produced electrochemically.[5]

Curiously, perchlorate can be produced by lightning discharges in the presence of chloride. Perchlorate has been detected in rain and snow samples from Florida and Lubbock, Texas.[6]

Uses

  • The dominant use of perchlorates is as oxidizers in propellants for rockets and fireworks. Of particular value is ammonium perchlorate composite propellant as a component of solid rocket fuel. In a related but smaller application, perchlorates are used extensively within the pyrotechnics industry and in certain munitions and for the manufacture of matches.[4]
  • Perchlorate is used to control static electricity in food packaging. Sprayed onto containers it stops statically charged food from clinging to plastic or paper/cardboard surface.[7]
  • Niche uses include lithium perchlorate, which decomposes exothermically to produce oxygen, useful in oxygen "candles" on spacecraft, submarines, and in other situations where a reliable backup oxygen supply is needed.[8]
  • Potassium perchlorate has, in the past, been used therapeutically to treat hyperthyroidism resulting from Graves' disease. It impedes the accumulation of iodide in the thyroid, which blocks thyroid hormone production.[9]

Chemical properties

The perchlorate ion is the least reactive oxidizer of the generalized chlorates. Perchlorate contains chlorine in its highest oxidation number. A table of reduction potentials of the four chlorates shows that, contrary to expectation, perchlorate is the weakest oxidant among the four in water.[10]

Ion Acidic reaction E° (V) Neutral/basic reaction E° (V)
Hypochlorite 2 H+ + 2 HOCl + 2 e → Cl2(g) + 2 H2O 1.63 ClO + H2O + 2 e → Cl + 2OH 0.89
Chlorite 6 H+ + 2 HOClO + 6 e → Cl2(g) + 4 H2O 1.64 ClO
2
+ 2 H2O + 4 e → Cl + 4 OH
0.78
Chlorate 12 H+ + 2 ClO
3
+ 10 e → Cl2(g) + 6 H2O
1.47 ClO
3
+ 3 H2O + 6 e → Cl + 6 OH
0.63
Perchlorate 16 H+ + 2 ClO
4
+ 14 e → Cl2(g) + 8 H2O
1.42 ClO
4
+ 4 H2O + 8 e → Cl + 8 OH
0.56

These data show that the perchlorate and chlorate are stronger oxidizers in acidic conditions than in basic conditions.

Gas phase measurements of heats of reaction (which allow computation of ΔHf°) of various chlorine oxides do follow the expected trend wherein Cl2O7 exhibits the largest endothermic value of ΔHf° (238.1 kJ/mol) while Cl2O exhibits the lowest endothermic value of ΔHf° (80.3 kJ/mol).[11]

The chlorine in the perchlorate anion is a closed shell atom and is well protected by the four oxygens. Most perchlorate compounds, especially salts of electropositive metals such as sodium perchlorate or potassium perchlorate, do not oxidize organic compounds until the mixture is heated. This property is useful in many applications, such as flares, where ignition is required to initiate a reaction. Ammonium perchlorate is stable when pure but can form potentially explosive mixtures with reactive metals or organic compounds. The PEPCON disaster destroyed a production plant for ammonium perchlorate when a fire caused the ammonium perchlorate stored on site to react with the aluminum that the storage tanks were constructed with and explode.

Potassium perchlorate has the lowest solubility of any alkali metal perchlorate (1.5 g in 100 ml of water at 25 °C).

Biology

Over 40 phylogenetically and metabolically diverse microorganisms capable of growth via perchlorate reduction[12] have been isolated since 1996. Most originate from the Proteobacteria but others include the Firmicutes, Moorella perchloratireducens and Sporomusa sp., and the archaeon Archaeoglobus fulgidus.[13][14] With the exception of A. fulgidus, all known microbes that grow via perchlorate reduction utilize the enzymes perchlorate reductase and chlorite dismutase, which collectively take perchlorate to innocuous chloride.[13] In the process, free oxygen (O2) is generated.[13]

Oxyanions of chlorine

Chlorine can assume oxidation states of −1, +1, +3, +5, or +7, an additional oxidation state of +4 is seen in the neutral compound chlorine dioxide ClO2, which has a similar structure. Several other chlorine oxides are also known.

Chlorine oxidation state −1 +1 +3 +5 +7
Name chloride hypochlorite chlorite chlorate perchlorate
Formula Cl ClO ClO
2
ClO
3
ClO
4
Structure The chloride ion The hypochlorite ion The chlorite ion The chlorate ion The perchlorate ion

Natural abundance

Terrestrial abundance

Naturally occurring perchlorate at its most abundant can be found comingled with deposits of sodium nitrate in the Atacama Desert of northern Chile. These deposits have been heavily mined as sources for nitrate-based fertilizers. Chilean nitrate is in fact estimated to be the source of around 81,000 tonnes (89,000 tons) of perchlorate imported to the U.S. (1909–1997). Results from surveys of ground water, ice, and relatively unperturbed deserts have been used to estimate a 100,000 to 3,000,000 tonnes (110,000 to 3,310,000 tons) "global inventory" of natural perchlorate presently on Earth.[15]

On Mars

In May 2008, the Wet Chemistry Laboratory (WCL) on board the 2007 Phoenix Mars lander performed the first wet chemical analysis of Martian soil. The analyses on three samples, two from the surface and one from a depth of 5 cm (2.0 in), revealed a slightly alkaline soil and low levels of salts typically found on Earth. Unexpected though was the presence of ~0.6% by weight perchlorate (ClO
4
), most likely as a mixture of 60% Ca(ClO4)2 and 40% Mg(ClO4)2.[16][17][18] These salts, formed from perchlorates, act as antifreeze and substantially lower the freezing point of water. Based on the temperature and pressure conditions on present-day Mars at the Phoenix lander site, conditions would allow a perchlorate salt solution to be stable in liquid form for a few hours each day during the summer.[19]

The possibility that the perchlorate was a contaminant brought from Earth has been eliminated by several lines of evidence. The Phoenix retro-rockets used ultra pure hydrazine and launch propellants consisting of ammonium perchlorate. Sensors on board Phoenix found no traces of ammonium, and thus the perchlorate in the quantities present in all three soil samples is indigenous to the Martian soil.

In 2006, a mechanism was proposed for the formation of perchlorates that is particularly relevant to the discovery of perchlorate at the Phoenix lander site. It was shown that soils with high concentrations of chloride converted to perchlorate in the presence of titanium dioxide and sunlight/ultraviolet light. The conversion was reproduced in the lab using chloride-rich soils from Death Valley.[20] Other experiments have demonstrated that the formation of perchlorate is associated with wide band gap semiconducting oxides.[21] In 2014, it was shown that perchlorate and chlorate can be produced from chloride minerals under Martian conditions via UV using only NaCl and silicate.[22]

Further findings of perchlorate and chlorate in the Martian meteorite EETA79001 [23] and by the Mars Curiosity rover in 2012-2013 support the notion that perchlorates are globally distributed throughout the Martian surface.[24][25][26] With concentrations approaching 0.5% and exceeding toxic levels on Martian soil, Martian perchlorates would present a serious challenge to human settlement,[27] as well as microorganisms.[28]

On September 28, 2015, NASA announced that analyses of spectral data from the Compact Reconnaissance Imaging Spectrometer for Mars instrument (CRISM) on board the Mars Reconnaissance Orbiter from four different locations where recurring slope lineae (RSL) are present found evidence for hydrated salts. The hydrated salts most consistent with the spectral absorption features are magnesium perchlorate, magnesium chlorate and sodium perchlorate. The findings strongly support the hypothesis that RSL form as a result of contemporary water activity on Mars.[29][30][31][32][33]

Contamination in environment

Perchlorate is of concern because of uncertainties about toxicity and health effects at low levels in drinking water, impact on ecosystems, and indirect exposure pathways for humans due to accumulation in vegetables.[9] Perchlorate is water-soluble, exceedingly mobile in aqueous systems, and can persist for many decades under typical groundwater and surface water conditions.[34] Detected perchlorate originates from disinfectants, bleaching agents, herbicides, and mostly from rocket propellants. Perchlorate is a byproduct of the production of a rocket fuel and fireworks.[3] The removal and recovery of the perchlorate compounds in explosives and rocket propellants include high-pressure water washout, which generate aqueous ammonium perchlorate.

In U.S. drinking water

Low levels of perchlorate have been detected in both drinking water and groundwater in 26 states in the U.S., according to the Environmental Protection Agency (EPA).[35] The chemical has been detected at levels as high as 5 µg/L at Joint Base Cape Cod (formerly Massachusetts Military Reservation), well over the Massachusetts state regulation of 2 µg/L.[36][37] Fireworks are also a source of perchlorate in lakes.[38]

At the Olin Flare Facility, Morgan Hill, California perchlorate contamination beneath the former flare manufacturing plant was first discovered in 2000, several years after the plant had closed. The plant had used potassium perchlorate as one of the ingredients during its 40 years of operation. By late 2003, the State of California and the Santa Clara Valley Water District had confirmed a groundwater plume currently extending over nine miles through residential and agricultural communities. The California Regional Water Quality Control Board and the Santa Clara Valley Water District have engaged in a major outreach effort, a water well testing program has been underway for about 1,200 residential, municipal, and agricultural wells. Large ion exchange treatment units are operating in three public water supply systems which include seven municipal wells with perchlorate detection. The potentially responsible parties, Olin Corporation and Standard Fuse Incorporated, have been supplying bottled water to nearly 800 households with private wells, and the Regional Water Quality Control Board has been overseeing cleanup efforts.[39]

The source of perchlorate in California was mainly attributed to two manufacturers in the southeast portion of the Las Vegas Valley in Nevada, where perchlorate has been produced for industrial use.[40] This led to perchlorate release into Lake Mead in Nevada and the Colorado River which affected regions of Nevada, California and Arizona, where water from this reservoir is used for consumption, irrigation and recreation for approximate half the population of these states.[3] Lake Mead has been attributed as the source of 90% of the perchlorate in Southern Nevada's drinking water. Based on sampling, perchlorate has been affecting 20 million people, with highest detection in Texas, southern California, New Jersey, and Massachusetts, but intensive sampling of the Great Plains and other middle state regions may lead to revised estimates with additional affected regions.[3] An action level of 18 μg/L has been adopted by several affected states.[34]

In food

In 2004, the chemical was found in cow's milk in California at an average level of 1.3 parts per billion (ppb, or µg/L), which may have entered the cows through feeding on crops exposed to water containing perchlorates.[41] A 2005 study suggested human breast milk had an average of 10.5 µg/L of perchlorate.[42]

In minerals and other natural occurrences

In some places, there is no clear source of perchlorate, and it may be naturally occurring. Natural perchlorate on earth was first identified in terrestrial nitrate deposits of the Atacama Desert in Chile as early as the 1880s[43] and for a long time considered a unique perchlorate source. The perchlorate released from historic use of Chilean nitrate based fertilizer which the U.S.imported by the hundreds of tons in the early 19th century can still be found in some groundwater sources of the United States.[44] Recent improvements in analytical sensitivity using ion chromatography based techniques have revealed a more widespread presence of natural perchlorate, particularly in subsoils of Southwest USA,[45] salt evaporites in California and Nevada,[46] Pleistocene groundwater in New Mexico,[47] and even present in extremely remote places such as Antarctica.[48] The data from these studies and others indicate that natural perchlorate is globally deposited on Earth with the subsequent accumulation and transport governed by the local hydrologic conditions.

Despite its importance to environmental contamination, the specific source and processes involved in natural perchlorate production remain poorly understood. Laboratory experiments in conjunction with isotopic studies[49] have implied that perchlorate may be produced on earth by oxidation of chlorine species through pathways involving ozone or its photochemical products.[50] Other studies have suggested that perchlorate can also be created by lightning activated oxidation of chloride aerosols (e.g., chloride in sea salt sprays),[51] and ultraviolet or thermal oxidation of chlorine (e.g., bleach solutions used in swimming pools) in water.[52][53][54]

From fertilizers

Although perchlorate as an environmental contaminant is usually associated with the storage, manufacture, and testing of solid rocket motors,[55] contamination of perchlorate has been focused in the use of fertilizer and its perchlorate release into ground water. Fertilizer leaves perchlorate anions to leak into the ground water and threaten the water supplies of many regions in the US.[55] One of the main sources of perchlorate contamination from fertilizer use was found to come from the fertilizer derived from Chilean caliche (calcium carbonate), because Chile has rich source of naturally occurring perchlorate anion.[56] Perchlorate in the solid fertilizer ranged from 0.7 to 2.0 mg g−1, variation of less than a factor of 3 and it is estimated that sodium nitrate fertilizers derived from Chilean caliche contain approximately 0.5–2 mg g−1 of perchlorate anion.[56] The direct ecological effect of perchlorate is not well known; its impact can be influenced by factors including rainfall and irrigation, dilution, natural attenuation, soil adsorption, and bioavailability.[56] Quantification of perchlorate concentrations in fertilizer components via ion chromatography revealed that in horticultural fertilizer components contained perchlorate ranging between 0.1 and 0.46%.[34] Perchlorate concentration was the highest in Chilean nitrate, ranging from 3.3 to 3.98%.[34]

Cleanup

There have been many attempts to eliminate perchlorate contamination. Current remediation technologies for perchlorate have downsides of high costs and difficulty in operation.[57] Thus, there have been interests in developing systems that would offer economic and green alternatives.[57]

Treatment ex situ and in situ

Several technologies can remove perchlorate, via treatments ex situ and in situ.

Ex situ treatments include ion exchange using perchlorate-selective or nitrite-specific resins, bioremediation using packed-bed or fluidized-bed bioreactors, and membrane technologies via electrodialysis and reverse osmosis.[58] In ex situ treatment via ion exchange, contaminants are attracted and adhere to the ion exchange resin because such resins and ions of contaminants have opposite charge.[59] As the ion of the contaminant adheres to the resin, another charged ion is expelled into the water being treated, in which then ion is exchanged for the contaminant.[59] Ion exchange technology has advantages of being well-suitable for perchlorate treatment and high volume throughput but has a downside that it does not treat chlorinated solvents. In addition, ex situ technology of liquid phase carbon adsorption is employed, where granular activated carbon (GAC) is used to eliminate low levels of perchlorate and pretreatment may be required in arranging GAC for perchlorate elimination.[58]

In situ treatments, such as bioremediation via perchlorate-selective microbes and permeable reactive barrier, are also being used to treat perchlorate.[58] In situ bioremediation has advantages of minimal above-ground infrastructure and its ability to treat chlorinated solvents, perchlorate, nitrate, and RDX simultaneously. However, it has a downside that it may negatively affect secondary water quality. In situ technology of phytoremediation could also be utilized, even though perchlorate phytoremediation mechanism is not fully founded yet.[58]

Bioremediation using perchlorate-reducing bacteria, which reduce perchlorate ions to harmless chloride, has also been proposed.[60]

Health effects

Thyroid inhibition

Perchlorate is a potent competitive inhibitor of the thyroid sodium-iodide symporter.[61] Thus, it has been used to treat hyperthyroidism since the 1950s.[62] At very high doses (70,000–300,000 ppb) the administration of potassium perchlorate was considered the standard of care in the United States, and remains the approved pharmacologic intervention for many countries.

In large amounts perchlorate interferes with iodine uptake into the thyroid gland. In adults, the thyroid gland helps regulate the metabolism by releasing hormones, while in children, the thyroid helps in proper development. The NAS, in its 2005 report, Health Implications of Perchlorate Ingestion, emphasized that this effect, also known as Iodide Uptake Inhibition (IUI) is not an adverse health effect. However, in January 2008, California's Department of Toxic Substances Control stated that perchlorate is becoming a serious threat to human health and water resources.[63] In 2010, the EPA's Office of the Inspector General determined that the agency's own perchlorate reference dose of 24.5 parts per billion protects against all human biological effects from exposure. This finding was due to a significant shift in policy at the EPA in basing its risk assessment on non-adverse effects such as IUI instead of adverse effects. The Office of the Inspector General also found that because the EPA's perchlorate reference dose is conservative and protective of human health further reducing perchlorate exposure below the reference dose does not effectively lower risk.[64]

Perchlorate affects only thyroid hormone. Because it is neither stored nor metabolized, effects of perchlorate on the thyroid gland are reversible, though effects on brain development from lack of thyroid hormone in fetuses, newborns, and children are not.[65]

Toxic effects of perchlorate have been studied in a survey of industrial plant workers who had been exposed to perchlorate, compared to a control group of other industrial plant workers who had no known exposure to perchlorate. After undergoing multiple tests, workers exposed to perchlorate were found to have a significant systolic blood pressure rise compared to the workers who were not exposed to perchlorate, as well as a significant decreased thyroid function compared to the control workers.[66]

A study involving healthy adult volunteers determined that at levels above 0.007 milligrams per kilogram per day (mg/(kg·d)), perchlorate can temporarily inhibit the thyroid gland's ability to absorb iodine from the bloodstream ("iodide uptake inhibition", thus perchlorate is a known goitrogen).[67] The EPA converted this dose into a reference dose of 0.0007 mg/(kg·d) by dividing this level by the standard intraspecies uncertainty factor of 10. The agency then calculated a "drinking water equivalent level" of 24.5 ppb by assuming a person weighs 70 kg (150 lb) and consumes 2 L (0.44 imp gal; 0.53 US gal) of drinking water per day over a lifetime.[68]

In 2006, a study reported a statistical association between environmental levels of perchlorate and changes in thyroid hormones of women with low iodine. The study authors were careful to point out that hormone levels in all the study subjects remained within normal ranges. The authors also indicated that they did not originally normalize their findings for creatinine, which would have essentially accounted for fluctuations in the concentrations of one-time urine samples like those used in this study.[69] When the Blount research was re-analyzed with the creatinine adjustment made, the study population limited to women of reproductive age, and results not shown in the original analysis, any remaining association between the results and perchlorate intake disappeared.[70] Soon after the revised Blount Study was released, Robert Utiger, a doctor with the Harvard Institute of Medicine, testified before the US Congress and stated: "I continue to believe that that reference dose, 0.007 milligrams per kilo (24.5 ppb), which includes a factor of 10 to protect those who might be more vulnerable, is quite adequate."[71]

At a 2013 presentation of a previously unpublished study, it was suggested that environmental exposure to perchlorate in pregnant women with hypothyroidism may be associated with significant risk of low IQ in their children.[72]

Lung toxicity

Some studies suggest that perchlorate has pulmonary toxic effects as well. Studies have been performed on rabbits where perchlorate has been injected into the trachea. The lung tissue was removed and analyzed, and it was found that perchlorate injected lung tissue showed several adverse effects when compared to the control group that had been intratracheally injected with saline. Adverse effects included inflammatory infiltrates, alveolar collapse, subpleural thickening, and lymphocyte proliferation.[73]

Aplastic anemia

In the early 1960s, potassium perchlorate use to treat Graves disease was implicated in the development of aplastic anemia—a condition where the bone marrow fails to produce new blood cells in sufficient quantity—in thirteen patients, seven of whom died.[74] Subsequent investigations have indicated the connection between administration of potassium perchlorate and development of aplastic anemia to be "equivocable at best", which means that the benefit of treatment, if it is the only known treatment, outweighs the risk, and it appeared a contaminant poisoned the 13.[75]

Regulation in the U.S.

Water

In 1998, perchlorate was included in the EPA Contaminant Candidate List, primarily due to its detection in California drinking water.[76][3]

In 2003, a federal district court in California found that the Comprehensive Environmental Response, Compensation and Liability Act applied, because perchlorate is ignitable, and therefore was a "characteristic" hazardous waste.[77]

In 2003, California's legislature enacted AB 826, the Perchlorate Contamination Prevention Act of 2003, requiring California's Department of Toxic Substances Control (DTSC) to adopt regulations specifying best management practices for perchlorate and perchlorate-containing substances. On December 31, 2005, the "Perchlorate Best Management Practices" were adopted and became operative on July 1, 2006.[78]

In early 2006, EPA issued a "Cleanup Guidance" and recommended a Drinking Water Equivalent Level (DWEL) for perchlorate of 24.5 µg/L. Both DWEL and Cleanup Guidance were based on a 2005 review of the existing research by the National Academy of Science (NAS).[79]

Lacking a federal standard, several states in the U.S. subsequently enacted their own drinking water standard for perchlorate including Massachusetts in 2006 and California in 2007. Other states, including Arizona, Maryland, Nevada, New Mexico, New York, and Texas have established non-enforceable, advisory levels for perchlorate.

It was not until 2008, that EPA issued an interim drinking water health advisory for perchlorate and with it a guidance and analysis concerning the impacts on the environment and drinking water.[80] California also issued guidance regarding perchlorate use.[81] Both the Department of Defense and some environmental groups voiced questions about the NAS report, but no credible science has emerged to challenge the NAS findings.

In February 2008, U.S. Food and Drug Administration reported that U.S. toddlers on average are being exposed to more than half of EPA's safe dose from food alone.[82] In March 2009, a Centers for Disease Control study found 15 brands of infant formula contaminated with perchlorate. Combined with existing perchlorate drinking water contamination, infants could be at risk for perchlorate exposure above the levels considered safe by EPA.

On February 11, 2011, EPA determined that perchlorate meets the Safe Drinking Water Act criteria for regulation as a contaminant.[80][83] The agency found that perchlorate may have an adverse effect on the health of persons and is known to occur in public water systems with a frequency and at levels that it presents a public health concern. Since then EPA has continued to determine what level of contamination is appropriate. The EPA prepared extensive responses to submitted public comments.[84]

In 2016, the Natural Resources Defense Council filed a lawsuit to accelerate EPA's regulation of perchlorate. A federal district court in New York issued a consent decree that initially required EPA to issue a proposed rule in October 2018, and a final rule in December 2019.[85] The modified court order requires EPA to issue a proposed rule by May 28, 2019.[86]

Other

FDA approved perchlorate use in food packaging in 2005.

References

  1. ^ "Perchlorate - PubChem Public Chemical Database". The PubChem Project. USA: National Center for Biotechnology Information.
  2. ^ Draft Toxicological Profile for Perchlorates, Agency for Toxic Substances and Disease Registry, U.S. Department of Health and Human Services, September, 2005.
  3. ^ a b c d e Kucharzyk, Katarzyna (2009). "Development of drinking water standards for perchlorate in the United States". Journal of Environmental Management. 91 (2): 303–310. doi:10.1016/j.jenvman.2009.09.023.
  4. ^ a b Helmut Vogt, Jan Balej, John E. Bennett, Peter Wintzer, Saeed Akbar Sheikh, Patrizio Gallone "Chlorine Oxides and Chlorine Oxygen Acids" in Ullmann's Encyclopedia of Industrial Chemistry 2002, Wiley-VCH. doi:10.1002/14356007.a06_483
  5. ^ Dotson R.L. (1993). "A novel electrochemical process for the production of ammonium perchlorate". Journal of Applied Electrochemistry. 23 (9): 897–904. doi:10.1007/BF00251024.
  6. ^ Kathleen Sellers, Katherine Weeks, William R. Alsop, Stephen R. Clough, Marilyn Hoyt, Barbara Pugh, Joseph Robb. Perchlorate: Environmental Problems and Solutions, 2007, p 9. Taylor & Francis Group, LLC.
  7. ^ McMullen Jenica, Ghassabian Akhgar, Kohn Brenda, Trasande Leonardo (2017). "Identifying Subpopulations Vulnerable to the Thyroid-Blocking Effects of Perchlorate and Thiocyanate". The Journal of Clinical Endocrinology & Metabolism. 102 (7): 2637–2645. doi:10.1210/jc.2017-00046.CS1 maint: Multiple names: authors list (link)
  8. ^ Markowitz, M. M.; Boryta, D. A.; Stewart, Harvey (1964). "Lithium Perchlorate Oxygen Candle. Pyrochemical Source of Pure Oxygen". Industrial & Engineering Chemistry Product Research and Development. 3 (4): 321–330. doi:10.1021/i360012a016.
  9. ^ a b Susarla Sridhar; Collette C. W.; Garrison A. W.; Wolfe N. L.; McCutcheon S. C. (1999). "Perchlorate Identification in Fertilizers". Environmental Science and Technology. 33 (19): 3469–3472. Bibcode:1999EnST...33.3469S. doi:10.1021/es990577k.
  10. ^ Cotton, F. Albert; Wilkinson, Geoffrey (1988), Advanced Inorganic Chemistry (5th ed.), New York: Wiley-Interscience, p. 564, ISBN 0-471-84997-9
  11. ^ Wagman, D. D.; Evans, W. H.; Parker, V. P.; Schumm, R. H.; Halow, I.; Bailey, S. M.; Churney, K. L.; Nuttall, R. L. J. Phys. Chem. Ref. Data Vol. 11(2); &169;1982 by the American Chemical Society and the American Institute of Physics.
  12. ^ J. Cameron Thrash, Jarrod Pollock, Tamas Torok, and John D. Coates, "Description of the novel perchlorate-reducing bacteria Dechlorobacter hydrogenophilus gen. nov., sp. nov. and Propionivibrio militaris, sp. nov.", Appl Microbiol Biotechnol. 2010 Mar; 86(1): 335–343. Published online 2009 Nov 18. doi: 10.1007/s00253-009-2336-6 PMCID: PMC2822220 PMID 19921177. Retrieved 12 April 2019.
  13. ^ a b c John D. Coates; Laurie A. Achenbach (2004). "Microbial perchlorate reduction: rocket-fuelled metabolism". Nature Reviews Microbiology. 2 (7): 569–580. doi:10.1038/nrmicro926. PMID 15197392.
  14. ^ Martin G. Liebensteiner, Martijn W. H. Pinkse, Peter J. Schaap, Alfons J. M. Stams, Bart P. Lomans (5 April 2013). "Archaeal (Per)Chlorate Reduction at High Temperature: An Interplay of Biotic and Abiotic Reactions". Science. 340 (6128): 85–87. Bibcode:2013Sci...340...85L. doi:10.1126/science.1233957. PMID 23559251.CS1 maint: Multiple names: authors list (link)
  15. ^ DuBois, Jennifer L.; Ojha, Sunil (2015). "Chapter 3, Section 2.2 Natural Abundance of Perchlorate on Earth". In Peter M.H. Kroneck and Martha E. Sosa Torres (ed.). Sustaining Life on Planet Earth: Metalloenzymes Mastering Dioxygen and Other Chewy Gases. Metal Ions in Life Sciences. 15. Springer. pp. 45–87. doi:10.1007/978-3-319-12415-5_3. ISBN 978-3-319-12414-8. PMC 5012666. PMID 25707466.
  16. ^ Hecht, M. H., S. P. Kounaves, R. Quinn; et al. (2009). "Detection of Perchlorate & the Soluble Chemistry of Martian Soil at the Phoenix Mars Lander Site". Science. 325 (5936): 64–67. Bibcode:2009Sci...325...64H. doi:10.1126/science.1172466. PMID 19574385.CS1 maint: Multiple names: authors list (link)
  17. ^ Kounaves S. P.; et al. (2010). "Wet Chemistry Experiments on the 2007 Phoenix Mars Scout Lander: Data Analysis and Results". J. Geophys. Res. 115 (E3): E00E10. Bibcode:2009JGRE..114.0A19K. doi:10.1029/2008JE003084.
  18. ^ Kounaves S. P.; et al. (2014). "Identification of the Perchlorate Parent Salts at the Phoenix Mars Landing Site and Possible Implications". Icarus. 232: 226–231. Bibcode:2014Icar..232..226K. doi:10.1016/j.icarus.2014.01.016.
  19. ^ Chevrier, V. C., Hanley, J., and Altheide, T.S. (2009). "Stability of perchlorate hydrates and their liquid solutions at the Phoenix landing site, Mars". Geophysical Research Letters. 36 (10): L10202. Bibcode:2009GeoRL..3610202C. doi:10.1029/2009GL037497.CS1 maint: Multiple names: authors list (link)
  20. ^ Miller, Glen. "Photooxidation of chloride to perchlorate in the presence of desert soils and titanium dioxide". American Chemical Society. March 29, 2006
  21. ^ Schuttlefield Jennifer D.; Sambur Justin B.; Gelwicks Melissa; Eggleston Carrick M.; Parkinson B. A. (2011). "Photooxidation of Chloride by Oxide Minerals: Implications for Perchlorate on Mars". J. Am. Chem. Soc. 133 (44): 17521–17523. doi:10.1021/ja2064878. PMID 21961793.
  22. ^ Carrier B. L.; Kounaves S. P. (2015). "The Origin of Perchlorates in the Martian Soil". Geophys. Res. Lett. 42 (10): 3746–3754. Bibcode:2015GeoRL..42.3739C. doi:10.1002/2015GL064290.
  23. ^ Kounaves S. P.; Carrier B. L.; O'Neil G. D.; Stroble S. T. & Clair M. W. (2014). "Evidence of Martian Perchlorate, Chlorate, and Nitrate in Mars Meteorite EETA79001: Implications for Oxidants and Organics". Icarus. 229: 206–213. Bibcode:2014Icar..229..206K. doi:10.1016/j.icarus.2013.11.012.
  24. ^ Adam Mann. "Look What We Found on Mars - Curiosity Rover Serves Up Awesome Science". Slate (magazine). 26 September 2013.
  25. ^ Chang, Kenneth (1 October 2013). "Hitting Pay Dirt on Mars". New York Times. Retrieved 2 October 2013.
  26. ^ Kerr Richard A (2013). "Pesky Perchlorates All Over Mars". Science. 340 (6129): 138. doi:10.1126/science.340.6129.138-b. PMID 23580505.
  27. ^ David, Leonard (June 13, 2013). "Toxic Mars: Astronauts Must Deal with Perchlorate on the Red Planet". Space.com. Retrieved May 9, 2017.
  28. ^ Mars covered in toxic chemicals that can wipe out living organisms, tests reveal. Ian Sample, The Guardian. 6 July 2017.
  29. ^ Webster, Guy; Agle, DC; Brown, Dwayne; Cantillo, Laurie (28 September 2015). "NASA Confirms Evidence That Liquid Water Flows on Today's Mars". Retrieved 28 September 2015.
  30. ^ Chang, Kenneth (28 September 2015). "NASA Says Signs of Liquid Water Flowing on Mars". New York Times. Retrieved 28 September 2015.
  31. ^ Ojha, Lujendra; Wilhelm, Mary Beth; Murchie, scortt L.; McEwen, Alfred S.; Wray, James J.; Hanley, Jennifer; Massé, Marion; Chojnacki, Matt (28 September 2015). "Spectral evidence for hydrated salts in recurring slope lineae on Mars". Nature Geoscience. 8 (11): 829–832. Bibcode:2015NatGe...8..829O. doi:10.1038/ngeo2546. Retrieved 28 September 2015.
  32. ^ Staff (28 September 2015). "Video Highlight (02:58) - NASA News Conference - Evidence of Liquid Water on Today's Mars". NASA. Retrieved 30 September 2015.
  33. ^ Staff (28 September 2015). "Video Complete (58:18) - NASA News Conference - Water Flowing on Present-Day Mars m". NASA. Retrieved 30 September 2015.
  34. ^ a b c d Susarla Sridhar; Collette T. W.; Garrison A. W.; Wolfe N. L.; McCutcheon S. C. (1999). "Perchlorate Identification in Fertilizers". Environmental Science and Technology. 33 (19): 3469–3472. Bibcode:1999EnST...33.3469S. doi:10.1021/es990577k.
  35. ^ Brandhuber, Philip; Clark, Sarah; Morley, Kevin (November 2009). "A review of perchlorate occurrence in public drinking water systems" (PDF). Journal American Water Works Association. 101 (11): 63–73.
  36. ^ Clausen, Jay (November 2001). "Perchlorate, Source and Distribution in Groundwater at Massachusetts Military Reservation" (PDF). Presentation at U.S. EPA Technical Support Project Semi-Annual Meeting, Cambridge, MA.
  37. ^ "Inorganic Chemical Maximum Contaminant Levels, Monitoring Requirements and Analytical Methods" (PDF). Massachusetts Office of Energy and Environmental Affairs. Code of Massachusetts Regulations (CMR), 310 CMR 22.06. Retrieved 2017-07-05.
  38. ^ "Fireworks Displays Linked To Perchlorate Contamination In Lakes". Science Daily. Rockville, MD. 2007-05-28.
  39. ^ "Perchlorate in the Pacific Southwest: California". EPA - Region 9. San Francisco, CA: EPA.
  40. ^ "Perchlorate". Las Vegas Valley Water District. Las Vegas, NV. Retrieved 2017-07-06.
  41. ^ Associated Press. "Toxic chemical found in California milk". MSNBC. June 22, 2004.
  42. ^ McKee, Maggie. "Perchlorate found in breast milk across US". New Scientist. February 23, 2005
  43. ^ Ericksen, G. E. "Geology and origin of the Chilean nitrate deposits"; U.S. Geological Survey Prof. Paper 1188; USGS: Reston, VA, 1981, 37 pp.
  44. ^ Böhlke J. K.; Hatzinger P. B.; Sturchio N. C.; Gu B.; Abbene I.; Mroczkowski S. J. (2009). "Atacama perchlorate as an agricultural contaminant in groundwater: Isotopic andchronologic evidence from Long Island, New York". Environmental Science & Technology. 43 (15): 5619–5625. Bibcode:2009EnST...43.5619B. doi:10.1021/es9006433.
  45. ^ Rao B.; Anderson T. A.; Orris G. J.; Rainwater K. A.; Rajagopalan S.; Sandvig R. M.; Scanlon B. R.; Stonestrom S. A.; Walvoord M. A.; Jackson W. A. (2007). "Widespread NaturalPerchlorate in Unsaturated zones of the Southwest United States". Environ. Sci. Technol. 41 (13): 4522–4528. Bibcode:2007EnST...41.4522R. doi:10.1021/es062853i.
  46. ^ Orris, G. J.; Harvey, G. J.; Tsui, D. T.; Eldridge, J. E. Preliminaryanalyses for perchlorate in selected natural materials and theirderivative products; USGS Open File Report 03-314; USGS, U.S.Government Printing Office: Washington, DC, 2003.
  47. ^ Plummer L. N.; Bohlke J. K.; Doughten M. W. (2005). "Perchlorate in Pleistocene and Holocene groundwater in North-Central New Mexico". Environ. Sci. Technol. 40 (6): 1757–1763. Bibcode:2006EnST...40.1757P. doi:10.1021/es051739h.
  48. ^ S. P. Kounaves; et al. (2010). "Natural Perchlorate in the Antarctic Dry Valleys and Implications for its Global Distribution and History". Environmental Science & Technology. 44 (7): 2360–2364. Bibcode:2010EnST...44.2360K. doi:10.1021/es9033606. PMID 20155929.
  49. ^ Böhlke, Karl John, Sturchio Neil C., Gu Baohua, Horita Juske, Brown Gilbert M., Jackson W. Andrew, Batista Jacimaria, Hatzinger Paul B. (2005). "Perchlorate isotope forensics". Analytical Chemistry. 77 (23): 7838–7842. doi:10.1021/ac051360d. PMID 16316196.CS1 maint: Multiple names: authors list (link)
  50. ^ Rao B., Anderson T. A., Redder A., Jackson W. A. (2010). "Perchlorate Formation by Ozone Oxidation of AqueousChlorine/Oxy-Chlorine Species: Role of ClxOy Radicals". Environ. Sci. Technol. 44 (8): 2961–2967. Bibcode:2010EnST...44.2961R. doi:10.1021/es903065f. PMID 20345093.CS1 maint: Multiple names: authors list (link)
  51. ^ Dasgupta P. K.; Martinelango P. K.; Jackson W. A.; Anderson T. A.; Tian K.; Tock R.W.; Rajagopalan S. (2005). "The origin of naturally occurring perchlorate: the role ofatmospheric processes". Environmental Science & Technology. 39 (6): 1569–1575. Bibcode:2005EnST...39.1569D. doi:10.1021/es048612x.
  52. ^ Rao B.; Estrada N; Mangold J.; Shelly M.; Gu B.; Jackson W. A. (2012). "Perchlorate production byphotodecomposition of aqueous chlorine". Environ. Sci. Technol. 46 (21): 11635–11643. Bibcode:2012EnST...4611635R. doi:10.1021/es3015277.
  53. ^ Stanford B. D.; Pisarenko A. N.; Snyder S. A.; Gordon G. (2011). "Perchlorate, bromate, and chlorate in hypochlorite solutions: Guidelines for utilities". Journal American Water Works Association. 103 (6): 71.
  54. ^ William E. Motzer (2001). "Perchlorate: Problems, Detection, and Solutions". Environmental Forensics. 2 (4): 301–311. doi:10.1006/enfo.2001.0059.
  55. ^ a b Magnuson Matthew L.; Urbansky Edward T.; Kelty Catherine A. (2000). "Determination of Perchlorate at Trace Levels in Drinking Water by Ion-Pair Extraction with Electrospray Ionization Mass Spectrometry". Analytical Chemistry. 72: 25–29. doi:10.1021/ac9909204.
  56. ^ a b c Urbansky T.; Brown S.K.; Magnuson M.L.; Kelty C.A. (2001). "Perchlorate levels in samples of sodium nitrate fertilizer derived from Chilean caliche". Environmental Pollution. 112 (3): 299–302. doi:10.1016/s0269-7491(00)00132-9.
  57. ^ a b "Eliminating Water Contamination by Inorganic Disinfection Byproducts". Hazen and Sawyer. Hazen and Sawyer.
  58. ^ a b c d "Technical Fact Sheet – Perchlorate" (PDF). US EPA. US EPA. 2013-04-23.
  59. ^ a b "ARA Perchlorate Contamination Solutions." Ion Exchange Perchlorate Treatment Solutions. ARA, n.d. Web. 25 Apr. 2014. <http://www.ara.com/perchlorate/Ion-Exchange-Perchlorate.html>.
  60. ^ Nirmala Bardiya and Jae-HoBae, "Dissimilatory perchlorate reduction: A review", Microbiological Research, Volume 166, Issue 4, 20 May 2011, pp. 237-254. Retrieved 12 April 2019.
  61. ^ Braverman, L. E.; He X.; Pino S.; et al. (2005). "The effect of perchlorate, thiocyanate, and nitrate on thyroid function in workers exposed to perchlorate long-term". J Clin Endocrinol Metab. 90 (2): 700–706. doi:10.1210/jc.2004-1821. PMID 15572417.
  62. ^ Godley, A. F.; Stanbury, J. B. (1954). "Preliminary experience in the treatment of hyperthyroidism with potassium perchlorate". J Clin Endocrinol Metab. 14 (1): 70–78. doi:10.1210/jcem-14-1-70. PMID 13130654.
  63. ^ "Perchlorate". California Department of Toxic Substances Control. Jan 26, 2008.
  64. ^ Scientific Analysis of Perchlorate: What We Found. Office of the Inspector General (Report). EPA. 19 April 2010.
  65. ^ J. Wolff (1998). "Perchlorate and the Thyroid Gland". Pharmacological Reviews. 50 (1): 89–105. PMID 9549759.
  66. ^ Chen HX, Shao YP, Wu FH, Li YP, Peng KL (Jan 2013). "[Health survey of plant workers for an occupational exposure to ammonium perchlorate]". Zhonghua Lao Dong Wei Sheng Zhi Ye Bing Za Zhi. 31 (1): 45–7. PMID 23433158.
  67. ^ Greer, M. A., Goodman, G., Pleuss, R. C., Greer, S. E. (2002). "Health effect assessment for environmental perchlorate contamination: The dose response for inhibition of thyroidal radioiodide uptake in humans" (free online). Environmental Health Perspectives. 110 (9): 927–937. doi:10.1289/ehp.02110927. PMC 1240994. PMID 12204829.CS1 maint: Multiple names: authors list (link)
  68. ^ "Perchlorate Guidance (Memorandum)" (PDF). EPA. January 26, 2006.
  69. ^ Benjamin C. Blount; James L. Pirkle; John D. Osterloh; Liza Valentin-Blasini & Kathleen L. Caldwell (2006). "Urinary Perchlorate and Thyroid Hormone Levels in Adolescent and Adult Men and Women Living in the United States". Environmental Health Perspectives. 114 (12): 1865–71. doi:10.1289/ehp.9466. PMC 1764147. PMID 17185277.
  70. ^ Tarone; et al. (2010). "The Epidemiology of Environmental Perchlorate Exposure and Thyroid Function: A Comprehensive Review". Journal of Occupational and Environmental Medicine. 52 (June): 653–60. doi:10.1097/JOM.0b013e3181e31955. PMID 20523234.
  71. ^ "Perchlorate: Health and Environmental Impacts of Unregulated Exposure". United States Congress. Retrieved 15 April 2012.
  72. ^ "Perchlorate Levels in Pregnancy Linked to Low Childhood IQ", by Nancy A. Melville, October 22, 2013
  73. ^ Wu F.; Chen H.; Zhou X.; Zhang R.; Ding M.; Liu Q.; Peng KL. (2013). "Pulmonary fibrosis effect of ammonium perchlorate exposure in rabbit". Arch Environ Occup Health. 68 (3): 161–5. doi:10.1080/19338244.2012.676105. PMID 23566323.
  74. ^ National Research Council (2005). "Perchlorate and the thyroid". Health implications of perchlorate ingestion. Washington, D.C: National Academies Press. p. 7. ISBN 978-0-309-09568-6. Retrieved on April 3, 2009 through Google Book Search.
  75. ^ Clark, J. J. J. (2000). "Toxicology of perchlorate". In Urbansky ET (ed.). Perchlorate in the environment. New York: Kluwer Academic/Plenum Publishers. pp. 19–20. ISBN 978-0-306-46389-1. Retrieved on April 3, 2009 through Google Book Search.
  76. ^ EPA (1998-03-02). "Announcement of the Drinking Water Contaminant Candidate List." Federal Register, 63 FR 10274
  77. ^ Castaic Lake Water Agency v. Whittaker, 272 F. Supp. 2d 1053, 1059–61 (C.D. Cal. 2003).
  78. ^ "Perchlorate". Managing Waste. Sacramento, CA: California Department of Toxic Substances Control. Retrieved 2017-05-28.
  79. ^ Committee to Assess the Health Implications of Perchlorate Ingestion, National Research Council (2005). Health Implications of Perchlorate Ingestion. Washington, DC: The National Academies Press. doi:10.17226/11202. ISBN 978-0-309-09568-6.
  80. ^ a b "Perchlorate in Drinking Water". Drinking Water Contaminants—Standards and Regulations. EPA. 2017-03-31.
  81. ^ "Perchlorate in Drinking Water". Drinking Water Systems. Sacramento, CA: California Department of Public Health. 2012-12-07. Archived from the original on 2013-02-06.
  82. ^ Renner, Rebecca (2008-03-15). "Perchlorate In Food". Environ. Sci. Technol. 42 (6): 1817. Bibcode:2008EnST...42.1817R. doi:10.1021/es0870552.
  83. ^ EPA (2011-02-11). "Drinking Water: Regulatory Determination on Perchlorate." 76 FR 7762
  84. ^ EPA-HQ-OW-2009-0297 "Docket ID" for EPA
  85. ^ Natural Resources Defense Council, Inc. v. United States Environmental Protection Agency and Gina McCarthy, 16 Civ. 1251 (ER). United States District Court for the Southern District of New York. Consent Decree filed October 17, 2016.
  86. ^ "Regulatory Update At-A-Glance". Washington, DC: Association of Metropolitan Water Agencies. Retrieved 2019-04-04.

External links

Ammonium perchlorate

Ammonium perchlorate ("AP") is an inorganic compound with the formula NH4ClO4. It is a colorless or white solid that is soluble in water. It is a powerful oxidizer. Combined with a fuel, it can be used as a rocket propellant. Its instability has involved it in a number of accidents, such as the PEPCON disaster.

Barium perchlorate

Barium perchlorate is a powerful oxidizing agent, with the formula Ba(ClO4)2. It is used in the pyrotechnic industry.

Barium perchlorate decomposes at 505 °C.

Caesium perchlorate

Caesium perchlorate or cesium perchlorate (CsClO4), is a perchlorate of caesium. It forms white crystals, which are sparingly soluble in cold water and ethanol. It dissolves more easily in hot water.

CsClO4 is the least soluble of the alkali metal perchlorates (followed by Rb, K, Li, and Na), a property which may be used for separatory purposes and even for gravimetric analysis. This low solubility played an important role in the characterization of francium as an alkali metal, as francium perchlorate coprecipitates with caesium perchlorate.

When heated, CsClO4 decomposes to caesium chloride above 250 °C. Like all perchlorates, it is a strong oxidant and may react violently with reducing agents and organic materials, especially at elevated temperatures.

Chlorine perchlorate

Chlorine perchlorate is the chemical compound with the formula Cl2O4. This chlorine oxide is an asymmetric oxide, with one chlorine atom in oxidation state +1 and the other +7, with proper formula ClOClO3. It is produced by the photolysis of chlorine dioxide at room temperature with 436 nm ultraviolet light:

2 ClO2 → ClOClO3Chlorine perchlorate can also be made the following reactions at −45 °C.

CsClO4 + ClOSO2F → Cs(SO3)F + ClOClO3Chlorine perchlorate is a pale greenish liquid which decomposes at room temperature.

Dichlorine hexoxide

Dichlorine hexoxide is the chemical compound with the molecular formula Cl2O6, which is correct for its gaseous state. However, in liquid or solid form, this chlorine oxide ionizes into the dark red ionic compound chloryl perchlorate [ClO2]+[ClO4]−, which may be thought of as the mixed anhydride of chloric and perchloric acids.

It is produced by reaction between chlorine dioxide and excess ozone:

2 ClO2 + 2 O3 → 2 ClO3 + 2 O2 → Cl2O6 + 2 O2

Flare

A flare, also sometimes called a fusee, is a type of pyrotechnic that produces a brilliant light or intense heat without an explosion. Flares are used for distress signalling, illumination, or defensive countermeasures in civilian and military applications. Flares may be ground pyrotechnics, projectile pyrotechnics, or parachute-suspended to provide maximum illumination time over a large area. Projectile pyrotechnics may be dropped from aircraft, fired from rocket or artillery, or deployed by flare guns or handheld percussive tubes.

Iodine-131

Iodine-131 (131I) is an important radioisotope of iodine discovered by Glenn Seaborg and John Livingood in 1938 at the University of California, Berkeley. It has a radioactive decay half-life of about eight days. It is associated with nuclear energy, medical diagnostic and treatment procedures, and natural gas production. It also plays a major role as a radioactive isotope present in nuclear fission products, and was a significant contributor to the health hazards from open-air atomic bomb testing in the 1950s, and from the Chernobyl disaster, as well as being a large fraction of the contamination hazard in the first weeks in the Fukushima nuclear crisis. This is because I-131 is a major fission product of uranium and plutonium, comprising nearly 3% of the total products of fission (by weight). See fission product yield for a comparison with other radioactive fission products. I-131 is also a major fission product of uranium-233, produced from thorium.

Due to its mode of beta decay, iodine-131 is notable for causing mutation and death in cells that it penetrates, and other cells up to several millimeters away. For this reason, high doses of the isotope are sometimes less dangerous than low doses, since they tend to kill thyroid tissues that would otherwise become cancerous as a result of the radiation. For example, children treated with moderate dose of I-131 for thyroid adenomas had a detectable increase in thyroid cancer, but children treated with a much higher dose did not. Likewise, most studies of very-high-dose I-131 for treatment of Graves disease have failed to find any increase in thyroid cancer, even though there is linear increase in thyroid cancer risk with I-131 absorption at moderate doses. Thus, iodine-131 is increasingly less employed in small doses in medical use (especially in children), but increasingly is used only in large and maximal treatment doses, as a way of killing targeted tissues. This is known as "therapeutic use".

Iodine-131 can be "seen" by nuclear medicine imaging techniques (i.e., gamma cameras) whenever it is given for therapeutic use, since about 10% of its energy and radiation dose is via gamma radiation. However, since the other 90% of radiation (beta radiation) causes tissue damage without contributing to any ability to see or "image" the isotope, other less-damaging radioisotopes of iodine such as iodine-123 (see isotopes of iodine) are preferred in situations when only nuclear imaging is required. The isotope I-131 is still occasionally used for purely diagnostic (i.e., imaging) work, due to its low expense compared to other iodine radioisotopes. Very small medical imaging doses of I-131 have not shown any increase in thyroid cancer. The low-cost availability of I-131, in turn, is due to the relative ease of creating I-131 by neutron bombardment of natural tellurium in a nuclear reactor, then separating I-131 out by various simple methods (i.e., heating to drive off the volatile iodine). By contrast, other iodine radioisotopes are usually created by far more expensive techniques, starting with reactor radiation of expensive capsules of pressurized xenon gas.

Iodine-131 is also one of the most commonly used gamma-emitting radioactive industrial tracer. Radioactive tracer isotopes are injected with hydraulic fracturing fluid to determine the injection profile and location of fractures created by hydraulic fracturing.Much smaller incidental doses of iodine-131 than those used in medical therapeutic procedures, are supposed by some studies to be the major cause of increased thyroid cancers after accidental nuclear contamination. These studies suppose that cancers happen from residual tissue radiation damage caused by the I-131, and should appear mostly years after exposure, long after the I-131 has decayed. Other studies can't find a correlation.

Lithium perchlorate

Lithium perchlorate is the inorganic compound with the formula LiClO4. This white or colourless crystalline salt is noteworthy for its high solubility in many solvents. It exists both in anhydrous form and as a trihydrate.

Magnesium perchlorate

Magnesium perchlorate is a powerful oxidizing agent, with the formula Mg(ClO4)2. It is also a superior drying agent for gas analysis.

Magnesium perchlorate decomposes at 250 °C. The heat of formation is -568.90 kJ mol−1.The enthalpy of solution is quite high, so reactions are done in large amounts of water to dilute it.

It is sold under the trade name anhydrone. Manufacture of this product on a semi-industrial scale was first performed by G. Frederick Smith in his garage in Urbana Illinois, but later at a permanent facility in Columbus, OH called G. Frederick Smith Chemical Co. He sold the magnesium perchlorate to A. H. Thomas Co., now Thomas Scientific, under the trade name Dehydrite.

It is used as desiccant to dry gas or air samples, but is no longer advised, for use as a general desiccant, due to hazards inherent in perchlorates. It is dried by heating at 220 °C under vacuum.

Magnesium perchlorate is created by the reaction of magnesium hydroxide and perchloric acid.

In 2011, a study conducted at the Georgia Institute of Technology revealed the presence of magnesium perchlorate on the planet Mars. Being a drying agent, magnesium perchlorate retains water from the atmosphere and may release it when conditions are favorable and temperature is above 250K. Because briny solutions that contain magnesium perchlorate have a lower melting point than that of pure water, their abundance on Mars could serve as evidence that liquid water may exist on its surface, where temperature and pressure conditions would ordinarily cause water to freeze.

Mare Boreum quadrangle

The Mare Boreum quadrangle is one of a series of 30 quadrangle maps of Mars used by the United States Geological Survey (USGS) Astrogeology Research Program. The Mare Boreum quadrangle is also referred to as MC-1 (Mars Chart-1). Its name derives from an older name for a feature that is now called Planum Boreum, a large plain surrounding the polar cap.The quadrangle covers all of the Martian surface north of latitude 65°. It includes the north polar ice cap, which has a swirl pattern and is roughly 1,100 km across. Mariner 9 in 1972 discovered a belt of sand dunes that ring the polar ice deposits, which is 500 km across in some places and may be the largest dune field in the solar system. The ice cap is surrounded by the vast plains of Planum Boreum and Vastitas Borealis. Close to the pole, there is a large valley, Chasma Boreale, that may have been formed from water melting from the ice cap. An alternative view is that it was made by winds coming off the cold pole. Another prominent feature is a smooth rise, formerly called Olympia Planitia. In the summer, a dark collar around the residual cap becomes visible; it is mostly caused by dunes. The quadrangle includes some very large craters that stand out in the north because the area is smooth with little change in topography. These large craters are Lomonosov and Korolev. Although smaller, the crater Stokes is also prominent.

The Phoenix lander landed on Vastitas Borealis within the Mare Boreum quadrangle at 68.218830° N and 234.250778° E on May 25, 2008.

The probe collected and analyzed soil samples in an effort to detect water and determine how hospitable the planet might once have been for life to grow. It remained active there until winter conditions became too harsh around five months later.After the mission ended the journal Science reported that chloride, bicarbonate, magnesium, sodium potassium, calcium, and possibly sulfate were detected in the samples analyzed by Phoenix. The pH was narrowed down to 7.7±0.5. Perchlorate (ClO4), a strong oxidizer at elevated temperatures, was detected. This was a significant discovery because the chemical has the potential of being used for rocket fuel and as a source of oxygen for future colonists. Also, under certain conditions perchlorate can inhibit life; however some microorganisms obtain energy from the substance (by anaerobic reduction). The chemical when mixed with water can greatly lower freezing points, in a manner similar to how salt is applied to roads to melt ice. So, perchlorate may be allowing small amounts of liquid water to form on Mars today. Gullies, which are common in certain areas of Mars, may have formed from perchlorate melting ice and causing water to erode soil on steep slopes.Much direct evidence was found for water at this location.

Nuclear fission product

Nuclear fission products are the atomic fragments left after a large atomic nucleus undergoes nuclear fission. Typically, a large nucleus like that of uranium fissions by splitting into two smaller nuclei, along with a few neutrons, the release of heat energy (kinetic energy of the nuclei), and gamma rays. The two smaller nuclei are the fission products. (See also Fission products (by element)).

About 0.2% to 0.4% of fissions are ternary fissions, producing a third light nucleus such as helium-4 (90%) or tritium (7%).

The fission products themselves are usually unstable and therefore radioactive; due to being relatively neutron-rich for their atomic number, many of them quickly undergo beta decay. This releases additional energy in the form of beta particles, antineutrinos, and gamma rays. Thus, fission events normally result in beta and gamma radiation, even though this radiation is not produced directly by the fission event itself.

The produced radionuclides have varying half-lives, and therefore vary in radioactivity. For instance, strontium-89 and strontium-90 are produced in similar quantities in fission, and each nucleus decays by beta emission. But 90Sr has a 30-year half-life, and 89Sr a 50.5-day half-life. Thus in the 50.5 days it takes half the 89Sr atoms to decay, emitting the same number of beta particles as there were decays, less than 0.4% of the 90Sr atoms have decayed, emitting only 0.4% of the betas. The radioactive emission rate is highest for the shortest lived radionuclides, although they also decay the fastest. Additionally, less stable fission products are less likely to decay to stable nuclides, instead decaying to other radionuclides, which undergo further decay and radiation emission, adding to the radiation output. It is these short lived fission products that are the immediate hazard of spent fuel, and the energy output of the radiation also generates significant heat which must be considered when storing spent fuel. As there are hundreds of different radionuclides created, the initial radioactivity level fades quickly as short lived radionuclides decay, but never ceases completely as longer lived radionuclides make up more and more of the remaining unstable atoms.

Perchloric acid

Perchloric acid is a mineral acid with the formula HClO4. Usually found as an aqueous solution, this colorless compound is a stronger acid than sulphuric acid and nitric acid. It is a powerful oxidizer when hot, but aqueous solutions up to approximately 70% by weight at room temperature are generally safe, only showing strong acid features and no oxidizing properties. Perchloric acid is useful for preparing perchlorate salts, especially ammonium perchlorate, an important rocket fuel component. Perchloric acid is dangerously corrosive and readily forms potentially explosive mixtures.

Potassium perchlorate

Potassium perchlorate is the inorganic salt with the chemical formula KClO4. Like other perchlorates, this salt is a strong oxidizer although it usually reacts very slowly with organic substances. This, usually obtained as a colorless, crystalline solid, is a common oxidizer used in fireworks, ammunition percussion caps, explosive primers, and is used variously in propellants, flash compositions, stars, and sparklers. It has been used as a solid rocket propellant, although in that application it has mostly been replaced by the higher performance ammonium perchlorate. KClO4 has the lowest solubility of the alkali metal perchlorates (1.5 g in 100 mL of water at 25 °C).

Pyrotechnic initiator

A pyrotechnic initiator (also initiator or igniter) is a device containing a pyrotechnic composition used primarily to ignite other, more difficult-to-ignite materials, e.g. thermites, gas generators, and solid-fuel rockets. The name is often used also for the compositions themselves.

Pyrotechnic initiators are often controlled electrically (called electro-pyrotechnic initiators), e.g. using a heated bridgewire or a bridge resistor. They are somewhat similar to blasting caps or other detonators, but they differ in that there is no intention to produce a shock wave. An example of such pyrotechnic initiator is an electric match.

Silver perchlorate

Silver perchlorate is the chemical compound with the formula AgClO4. This white solid forms a monohydrate and is mildly deliquescent. It is a useful source of the Ag+ ion, although the presence of perchlorate presents risks. It is used as a catalyst in organic chemistry.

Sodium perchlorate

Sodium perchlorate is the inorganic compound with the chemical formula NaClO4. It is a white crystalline, hygroscopic solid that is highly soluble in water and in alcohol. It usually encountered as the monohydrate. The compound is noteworthy as the most water-soluble of the common perchlorate salts.

Sodium perchlorate is present on the planet Mars.

Titanium perchlorate

Titanium perchlorate is a molecular compound of titanium and perchlorate groups with formula Ti(ClO4)4. Anhydrous titanium perchlorate decomposes explosively at 130°C and melts at 85°C with a slight decomposition. It can sublime in a vacuum as low as 70°C, and can form vapour at up to 120°. Titanium perchlorate is quite volatile. It has density 2.35. It decomposes to TiO2, ClO2 and dioxygen O2 Also TiO(ClO4)2 is formed during decomposition.Ti(ClO4)4 → TiO2 + 4ClO2 + 3O2 ΔH=+6 kcal/mol.

Vanadyl perchlorate

Vanadyl perchlorate or vanadyl triperchlorate is a golden yellow coloured liquid or crystalline compound of vanadium, oxygen and perchlorate group. The substance consists of molecules covalently bound and is quite volatile.

Compounds containing perchlorate group
HClO4 He
LiClO4 Be(ClO4)2 B(ClO
4
)
4

B(ClO4)3
ROClO3 N(ClO4)3
NH4ClO4
NOClO4
O FClO4 Ne
NaClO4 Mg(ClO4)2 Al(ClO4)3 Si P S ClO−4
ClOClO3
Cl2O7
Ar
KClO4 Ca(ClO4)2 Sc(ClO4)3 Ti(ClO4)4 VO(ClO4)3
VO2(ClO4)
Cr(ClO4)3 Mn(ClO4)2 Fe(ClO4)3 Co(ClO4)2,
Co(ClO4)3
Ni(ClO4)2 Cu(ClO4)2 Zn(ClO4)2 Ga(ClO4)3 Ge As Se Br Kr
RbClO4 Sr(ClO4)2 Y(ClO4)3 Zr(ClO4)4 Nb(ClO5)4 Mo Tc Ru Rh(ClO4)3 Pd(ClO4)2 AgClO4 Cd(ClO4)2 In(ClO4)3 Sn(ClO4)4 Sb TeO(ClO4)2 I Xe
CsClO4 Ba(ClO4)2   Hf(ClO4)4 Ta(ClO5)5 W Re Os Ir Pt Au Hg2(ClO4)2,
Hg(ClO4)2
Tl(ClO4)3 Pb(ClO4)2 Bi(ClO4)3 Po At Rn
FrClO4 Ra   Rf Db Sg Bh Hs Mt Ds Rg Cn Nh Fl Mc Lv Ts Og
La Ce(ClO4)x Pr Nd Pm Sm(ClO4)3 Eu(ClO4)3 Gd(ClO4)3 Tb(ClO4)3 Dy(ClO4)3 Ho(ClO4)3 Er(ClO4)3 Tm(ClO4)3 Yb(ClO4)3 Lu(ClO4)3
Ac Th(ClO4)4 Pa UO2(ClO4)2 Np Pu Am Cm Bk Cf Es Fm Md No Lr
Thyroid therapy (H03)
Thyroid hormones
Antithyroid preparations
Receptor
(ligands)
Transporter
(blockers)
Enzyme
(inhibitors)
Others

This page is based on a Wikipedia article written by authors (here).
Text is available under the CC BY-SA 3.0 license; additional terms may apply.
Images, videos and audio are available under their respective licenses.