Difference between revisions of "Manganate"
Physchim62 (talk | contribs) (→Preparation) |
Physchim62 (talk | contribs) (→Manganate(VI)) |
||
| Line 2: | Line 2: | ||
==Manganate(VI)== | ==Manganate(VI)== | ||
| − | The manganate(VI) ion is tetrahedral, similar to sulfate or chromate: indeed, manganates are often isostructural with sulfates and chromates. As a d<sup>1</sup> ion, it is [[Paramagnetism|paramagnetic]]. | + | The manganate(VI) ion is tetrahedral, similar to sulfate or chromate: indeed, manganates are often isostructural with sulfates and chromates. As a d<sup>1</sup> ion, it is [[Paramagnetism|paramagnetic]]. Manganates are dark green in colour, with a visible absorption maximum of ''λ''<sub>max</sub> = 606 nm (''ε'' = {{nowrap|1710 dm<sup>3</sup> mol<sup>−1</sup> cm<sup>−1</sup>}}). |
===Preparation=== | ===Preparation=== | ||
[[Sodium manganate|Sodium]] and [[potassium manganate]]s are usually prepared in the laboratory by stirring the equivalent [[permanganate]] in a concentrated solution (5–10 M) of the [[hydroxide]] for 24 hours.<ref>{{citation | journal = J. Chem. Soc. | year = 1956 | pages = 3373–80 | doi = 10.1039/JR9560003373 | title = Structure and reactivity of the oxy-anions of transition metals. Part I. The manganese oxy-anions | first1 = A. | last1 = Carrington | first2 = M. C. R. | last2 = Symons}}.</ref> | [[Sodium manganate|Sodium]] and [[potassium manganate]]s are usually prepared in the laboratory by stirring the equivalent [[permanganate]] in a concentrated solution (5–10 M) of the [[hydroxide]] for 24 hours.<ref>{{citation | journal = J. Chem. Soc. | year = 1956 | pages = 3373–80 | doi = 10.1039/JR9560003373 | title = Structure and reactivity of the oxy-anions of transition metals. Part I. The manganese oxy-anions | first1 = A. | last1 = Carrington | first2 = M. C. R. | last2 = Symons}}.</ref> | ||
| − | :{{nowrap|4 MnO{{su|b=4|p=−}}}} + {{nowrap|4 OH<sup>−</sup>}} → {{nowrap|4 MnO{{su|b=4|p=2−}} + {{nowrap|2 H<sub>2</sub>O}} + O<sub>2</sub> | + | :{{nowrap|4 MnO{{su|b=4|p=−}}}} + {{nowrap|4 OH<sup>−</sup>}} → {{nowrap|4 MnO{{su|b=4|p=2−}}}} + {{nowrap|2 H<sub>2</sub>O}} + O<sub>2</sub> |
Potassium manganate is prepared industrially, as an intermediate to [[potassium permanganate]], by dissolving [[manganese dioxide]] in molten [[potassium hydroxide]] with [[potassium nitrate]] or air as the [[oxidizing agent]].<ref name="C&W"/> | Potassium manganate is prepared industrially, as an intermediate to [[potassium permanganate]], by dissolving [[manganese dioxide]] in molten [[potassium hydroxide]] with [[potassium nitrate]] or air as the [[oxidizing agent]].<ref name="C&W"/> | ||
| Line 13: | Line 13: | ||
===Disproportionation=== | ===Disproportionation=== | ||
Manganates are unstable towards disproportionation in all but the most alkaline of [[aqueous solution]]s.<ref name="C&W"/> The ultimate products are [[permanganate]] and [[manganese dioxide]], but the [[Chemical kinetics|kinetics]] are complex and the mechanism may involved protonated and/or manganese(V) species.<ref>{{citation | title = Kinetics of the disproportionation of manganate in acid solution | first1 = Joan H. | last1 = Sutter | first2 = Kevin | last2 = Colquitt | first3 = John R. | last3 = Sutter | journal = Inorg. Chem. | year = 1974 | volume = 13 | issue = 6 | pages = 1444–46 | doi = 10.1021/ic50136a037}}.</ref> | Manganates are unstable towards disproportionation in all but the most alkaline of [[aqueous solution]]s.<ref name="C&W"/> The ultimate products are [[permanganate]] and [[manganese dioxide]], but the [[Chemical kinetics|kinetics]] are complex and the mechanism may involved protonated and/or manganese(V) species.<ref>{{citation | title = Kinetics of the disproportionation of manganate in acid solution | first1 = Joan H. | last1 = Sutter | first2 = Kevin | last2 = Colquitt | first3 = John R. | last3 = Sutter | journal = Inorg. Chem. | year = 1974 | volume = 13 | issue = 6 | pages = 1444–46 | doi = 10.1021/ic50136a037}}.</ref> | ||
| + | |||
| + | ===Manganic acid=== | ||
| + | '''Manganic acid''' cannot be formed because of its rapid disproportionation. However, its second [[acid dissociation constant]] has been estimated by [[pulse radiolysis]] techniques: | ||
| + | :HMnO{{su|b=4|p=−}} {{eqm}} MnO{{su|b=4|p=2−}} + H<sup>+</sup> p''K''<sub>a</sub> = 7.4 ± 0.1 | ||
| + | |||
| + | ==Manganate(V)== | ||
| + | The manganate(V) anion, MnO{{su|b=4|p=3−}}, known trivially as '''hypomanganate''' and systematically as '''tetraoxidomanganate(3−)''', is a bright blue species with a visible absorption maximum of ''λ''<sub>max</sub> = 670 nm (''ε'' = {{nowrap|900 dm<sup>3</sup> mol<sup>−1</sup> cm<sup>−1</sup>}}). It is unstable towards dispropotionation to manganate(VI) and [[manganese dioxide]], although the reaction is slow in very alkaline solution (''c''(OH<sup>−</sup>) = {{nowrap|5–10 mol dm<sup>−3</sup>}}). | ||
| + | |||
| + | Hypomanganates may be prepared by the careful reduction of manganates with [[sulfite]], [[hydrogen peroxide]] or [[Mandelic acid|mandelate]]. Only [[potassium hypomanganate]] has been studied to any significant extent. Hypomanganic acid cannot be formed because of its rapid disproportionation, but its third [[acid dissociation constant]] has been estimated by [[pulse radiolysis]] techniques: | ||
==References== | ==References== | ||
Revision as of 18:06, 26 June 2010
In inorganic nomenclature, a manganate is any negatively charged moelcular entity with manganese as the central atom. However, the name is usually used to refer to the tetraoxidomanganate(2−) anion, MnO2−4, also known as manganate(VI) because it contains manganese in the +6 oxidation state. Manganates are the only known manganese(VI) compounds.[1]
Contents
Manganate(VI)
The manganate(VI) ion is tetrahedral, similar to sulfate or chromate: indeed, manganates are often isostructural with sulfates and chromates. As a d1 ion, it is paramagnetic. Manganates are dark green in colour, with a visible absorption maximum of λmax = 606 nm (ε = 1710 dm3 mol−1 cm−1).
Preparation
Sodium and potassium manganates are usually prepared in the laboratory by stirring the equivalent permanganate in a concentrated solution (5–10 M) of the hydroxide for 24 hours.[2]
- 4 MnO−4 + 4 OH− → 4 MnO2−4 + 2 H2O + O2
Potassium manganate is prepared industrially, as an intermediate to potassium permanganate, by dissolving manganese dioxide in molten potassium hydroxide with potassium nitrate or air as the oxidizing agent.[1]
- 2 MnO2 + 4 OH− + O2 → 2 MnO2−4 + 2 H2O
Disproportionation
Manganates are unstable towards disproportionation in all but the most alkaline of aqueous solutions.[1] The ultimate products are permanganate and manganese dioxide, but the kinetics are complex and the mechanism may involved protonated and/or manganese(V) species.[3]
Manganic acid
Manganic acid cannot be formed because of its rapid disproportionation. However, its second acid dissociation constant has been estimated by pulse radiolysis techniques:
- HMnO−4 ⇌ MnO2−4 + H+ pKa = 7.4 ± 0.1
Manganate(V)
The manganate(V) anion, MnO3−4, known trivially as hypomanganate and systematically as tetraoxidomanganate(3−), is a bright blue species with a visible absorption maximum of λmax = 670 nm (ε = 900 dm3 mol−1 cm−1). It is unstable towards dispropotionation to manganate(VI) and manganese dioxide, although the reaction is slow in very alkaline solution (c(OH−) = 5–10 mol dm−3).
Hypomanganates may be prepared by the careful reduction of manganates with sulfite, hydrogen peroxide or mandelate. Only potassium hypomanganate has been studied to any significant extent. Hypomanganic acid cannot be formed because of its rapid disproportionation, but its third acid dissociation constant has been estimated by pulse radiolysis techniques:
References
- ↑ 1.0 1.1 1.2 Cotton, F. Albert; Wilkinson, Geoffrey Advanced Inorganic Chemistry, 4th ed.; Wiley: New York, 1980; p 746. ISBN 0-471-02775-8.
- ↑ Carrington, A.; Symons, M. C. R. Structure and reactivity of the oxy-anions of transition metals. Part I. The manganese oxy-anions. J. Chem. Soc. 1956, 3373–80. DOI: 10.1039/JR9560003373.
- ↑ Sutter, Joan H.; Colquitt, Kevin; Sutter, John R. Kinetics of the disproportionation of manganate in acid solution. Inorg. Chem. 1974, 13 (6), 1444–46. DOI: 10.1021/ic50136a037.
| Error creating thumbnail: Unable to save thumbnail to destination |
This page is currently licensed under the Creative Commons Attribution 3.0 Unported license and any later versions of that license. |
