Difference between revisions of "Chlorine oxoanions"
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'''Chlorine [[Oxoanion|oxoanions]]''' are negatively charged [[polyatomic ion]]s consisting of [[chlorine]] and [[oxygen]]. The chlorine atom is in an odd-number positive oxidation state, and the atom is surrounded by oxygen atoms and lone pairs of electrons. They are all oxidizing, but their stability varies, and some may have reducing tendencies. The four types of chlorine oxoanions are hypochlorites, chlorites, chlorates, and perchlorates. | '''Chlorine [[Oxoanion|oxoanions]]''' are negatively charged [[polyatomic ion]]s consisting of [[chlorine]] and [[oxygen]]. The chlorine atom is in an odd-number positive oxidation state, and the atom is surrounded by oxygen atoms and lone pairs of electrons. They are all oxidizing, but their stability varies, and some may have reducing tendencies. The four types of chlorine oxoanions are hypochlorites, chlorites, chlorates, and perchlorates. | ||
| + | |||
| + | ==Chlorine oxoanions== | ||
| + | *[[Hypochlorite]]: Formula ClO<sup>-</sup> | ||
| + | *[[Chlorite]]: Formula ClO<sub>2</sub><sup>-</sup> | ||
| + | *[[Chlorate]]: Formula ClO<sub>3</sub><sup>-</sup> | ||
| + | *[[Perchlorate]]: Formula ClO<sub>4</sub><sup>-</sup> | ||
==General properties== | ==General properties== | ||
| − | The oxoanions of chlorine are all quite unusual in their properties. As the oxygen content of the anion increases, the oxidizing ability of the ion actually decreases, as they become kinetically poorer oxidizers. More oxygenated anions are more stable with respect to oxidation and reduction. With light heating, hypochlorites will disproportionate into chlorates and chlorides, and chlorates will | + | The oxoanions of chlorine are all quite unusual in their properties. As the oxygen content of the anion increases, the oxidizing ability of the ion actually decreases, as they become kinetically poorer oxidizers. More oxygenated anions are more stable with respect to oxidation and reduction. With light heating, hypochlorites will disproportionate into chlorates and chlorides, and chlorates will disproportionate into perchlorates and chlorides. Perchlorates will convert to chlorides, with the loss of oxygen, in the presence of intense heat and or a reducing agent. |
==Availability== | ==Availability== | ||
| − | Hypochlorites are sold as bleach solutions ([[sodium hypochlorite]]) or powders ([[calcium hypochlorite]]). Chlorites are also available as industrial bleach. Chlorates and perchlorates are sometimes available as weed killers, though most have been phased out, due to their use in terrorist bombings. | + | Hypochlorites are sold as bleach solutions ([[sodium hypochlorite]]) or powders ([[calcium hypochlorite]]). Chlorites are also available as industrial bleach. Chlorates and perchlorates are sometimes available as weed killers, though most have been phased out, due to fire hazard and their use in terrorist bombings. |
==Preparation== | ==Preparation== | ||
| Line 19: | Line 25: | ||
Hypochlorites and chlorites will release chlorine and or chlorine dioxide in contact with a strong acid, which are higly toxic. They also tend to slowly decompose, yielding the above gaseous products as well as chlorates. Chlorates and perchlorates are strong oxidizers and must be kept away from any reducing agents and organic compounds. They are also toxic if ingested and may affect the thyroid gland. | Hypochlorites and chlorites will release chlorine and or chlorine dioxide in contact with a strong acid, which are higly toxic. They also tend to slowly decompose, yielding the above gaseous products as well as chlorates. Chlorates and perchlorates are strong oxidizers and must be kept away from any reducing agents and organic compounds. They are also toxic if ingested and may affect the thyroid gland. | ||
| − | ===Storage== | + | ===Storage=== |
| + | Hypochlorites and chlorites are unstable, and should be kept in plastic bottles. Chlorates and perchlorates are more stable, and can be kept in glass or plastic bottles. | ||
===Disposal=== | ===Disposal=== | ||
| + | Hypochlorites and chlorites can be neutralized with sodium or potassium sulfite, thiosulfate or metabisulfite. Chlorates can also be neutralized with acidified sulfites, metabisulfites or thiosulfates, although a mixture of sulfuric acid and ferrous ammonium sulfate can also be used. Perchlorates are a bit more complicated to safely neutralize. A strong reducing agent, like powdered iron added to an aqueous solution containing the perchlorate and exposed to strong UV light will reduce the perchlorate to chloride.<ref>[https://books.google.ro/books?id=gjfSBwAAQBAJ&pg=PA106&lpg=PA106&dq=perchlorate+neutralization&source=bl&ots=ztEPz18eE-&sig=oHHZzByFHPVjYIVxbvzGl4SC_xQ&hl=en&sa=X&ved=0ahUKEwi5gpTJm_XJAhVBURoKHb3LB4o4ChDoAQguMAc#v=onepage&q=perchlorate%20neutralization&f=false Perchlorate in the Environment (2000), Edward Todd Urbansky, pag. 106]</ref> Heating perchlorates with elemental iron at 200 °C for an hour destroyed 98% of the perchlorate.<ref>http://www.sciencedirect.com/science/article/pii/S0304389405005364</ref> | ||
==References== | ==References== | ||
| + | <references/> | ||
===Relevant Sciencemadness threads=== | ===Relevant Sciencemadness threads=== | ||
| + | *[http://www.sciencemadness.org/talk/viewthread.php?tid=7724 Determining the nomenclature of oxyanions] | ||
[[Category:Anions]] | [[Category:Anions]] | ||
[[Category:Oxoanions]] | [[Category:Oxoanions]] | ||
[[Category:Chlorine oxoanions]] | [[Category:Chlorine oxoanions]] | ||
Latest revision as of 18:35, 15 November 2019
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Chlorine oxoanions are negatively charged polyatomic ions consisting of chlorine and oxygen. The chlorine atom is in an odd-number positive oxidation state, and the atom is surrounded by oxygen atoms and lone pairs of electrons. They are all oxidizing, but their stability varies, and some may have reducing tendencies. The four types of chlorine oxoanions are hypochlorites, chlorites, chlorates, and perchlorates.
Contents
Chlorine oxoanions
- Hypochlorite: Formula ClO-
- Chlorite: Formula ClO2-
- Chlorate: Formula ClO3-
- Perchlorate: Formula ClO4-
General properties
The oxoanions of chlorine are all quite unusual in their properties. As the oxygen content of the anion increases, the oxidizing ability of the ion actually decreases, as they become kinetically poorer oxidizers. More oxygenated anions are more stable with respect to oxidation and reduction. With light heating, hypochlorites will disproportionate into chlorates and chlorides, and chlorates will disproportionate into perchlorates and chlorides. Perchlorates will convert to chlorides, with the loss of oxygen, in the presence of intense heat and or a reducing agent.
Availability
Hypochlorites are sold as bleach solutions (sodium hypochlorite) or powders (calcium hypochlorite). Chlorites are also available as industrial bleach. Chlorates and perchlorates are sometimes available as weed killers, though most have been phased out, due to fire hazard and their use in terrorist bombings.
Preparation
Chlorine oxoanion salts can be prepared by electrolysis of an alkali metal chloride and specific voltages.
Projects
- Flash powder
- Make chlorine
Handling
Safety
Hypochlorites and chlorites will release chlorine and or chlorine dioxide in contact with a strong acid, which are higly toxic. They also tend to slowly decompose, yielding the above gaseous products as well as chlorates. Chlorates and perchlorates are strong oxidizers and must be kept away from any reducing agents and organic compounds. They are also toxic if ingested and may affect the thyroid gland.
Storage
Hypochlorites and chlorites are unstable, and should be kept in plastic bottles. Chlorates and perchlorates are more stable, and can be kept in glass or plastic bottles.
Disposal
Hypochlorites and chlorites can be neutralized with sodium or potassium sulfite, thiosulfate or metabisulfite. Chlorates can also be neutralized with acidified sulfites, metabisulfites or thiosulfates, although a mixture of sulfuric acid and ferrous ammonium sulfate can also be used. Perchlorates are a bit more complicated to safely neutralize. A strong reducing agent, like powdered iron added to an aqueous solution containing the perchlorate and exposed to strong UV light will reduce the perchlorate to chloride.[1] Heating perchlorates with elemental iron at 200 °C for an hour destroyed 98% of the perchlorate.[2]
References
- ↑ Perchlorate in the Environment (2000), Edward Todd Urbansky, pag. 106
- ↑ http://www.sciencedirect.com/science/article/pii/S0304389405005364
