Difference between revisions of "Manganate"
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| − | In [[inorganic nomenclature]], a '''manganate''' is any negatively charged [[ | + | In [[inorganic nomenclature]], a '''manganate''' is any negatively charged [[molecular entity]] with [[manganese]] as the central atom. However, the name is usually used to refer to the '''tetraoxidomanganate(2−)''' anion, MnO{{su|b=4|p=2−}}, also known as '''manganate(VI)''' because it contains manganese in the +6 [[oxidation state]]. Manganates are the only known manganese(VI) compounds.<ref name="C&W">{{Cotton&Wilkinson4th|page=746}}.</ref> |
==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]]. 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>}}). | + | 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>}}).<ref name="JCS">{{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><ref name="Mandelate">{{citation | title = Reduction of manganate(VI) by mandelic acid and its significance for development of a general mechanism of oxidation of organic compounds by high-valent transition metal oxides | first1 = Donald G. | last1 = Lee | first2 = Tao | last2 = Chen | journal = J. Am. Chem. Soc. | year = 1993 | volume = 115 | issue = 24 | pages = 11231–36 | doi = 10.1021/ja00077a023}}.</ref> |
===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. | + | [[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. |
:{{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> | ||
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===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><ref name="Epot">{{citation | title = Rate of the MnO<sub>4</sub><sup>−</sup>/MnO<sub>4</sub><sup>2−</sup> and MnO<sub>4</sub><sup>2−</sup>/MnO<sub>4</sub><sup>3−</sup> electrode reactions in alkaline solutions at solid electrodes | first1 = K. | last1 = Sekula-Brzezińska | first2 = P. K. | last2 = Wrona | first3 = Z. | last3 = Galus | journal = Electrochim. Acta | year = 1979 | volume = 24 | issue = 5 | pages = 555–63 | doi = 10.1016/0013-4686(79)85032-X}}.</ref> |
===Manganic acid=== | ===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: | + | '''Manganic acid''' cannot be formed because of its rapid disproportionation. However, its second [[acid dissociation constant]] has been estimated by [[pulse radiolysis]] techniques:<ref name="PulseRad">{{citation | title = Studies of Manganate(V), -(VI), and -(VII) Tetraoxyanions by Pulse Radiolysis. Optical Spectra of Protonated Forms | first1 = J. D. | last1 = Rush | first2 = B. H. J. | last2 = Bielski | journal = Inorg. Chem. | year = 1995 | volume = 34 | issue = 23 | pages = 5832–38 | doi = 10.1021/ic00127a022}}.</ref> |
:HMnO{{su|b=4|p=−}} {{eqm}} MnO{{su|b=4|p=2−}} + H<sup>+</sup> p''K''<sub>a</sub> = 7.4 ± 0.1 | :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)== | ==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>}}). | + | The manganate(V) anion, MnO{{su|b=4|p=3−}}, known trivially as '''hypomanganate''' and systematically as '''tetraoxidomanganate(3−)''', is a bright blue species<ref name="G&E">{{Greenwood&Earnshaw1st|pages=1221–22}}.</ref> 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>}}).<ref name="JCS"/><ref name="Mandelate"/> 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>}}).<ref name="G&E"/> |
| − | 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: | + | Hypomanganates may be prepared by the careful reduction of manganates with [[sulfite]],<ref name="G&E"/> [[hydrogen peroxide]]<ref>{{citation | title = Oxidation of hydrocarbons. 18. Mechanism of the reaction between permanganate and carbon-carbon double bonds | first1 = Donald G. | last1 = Lee | first2 = Tao | last2 = Chen | journal = J. Am. Chem. Soc. | year = 1989 | volume = 111 | issue = 19 | pages = 7534–38 | doi = 10.1021/ja00201a039}}.</ref> or [[Mandelic acid|mandelate]].<ref name="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:<ref name="PulseRad"/> |
| + | :HMnO{{su|b=4|p=2−}} {{eqm}} MnO{{su|b=4|p=3−}} + H<sup>+</sup> p''K''<sub>a</sub> = {{nowrap|13.7 ± 0.2}} | ||
| + | |||
| + | ==Manganate(IV)== | ||
| + | The manganate(IV) anion has been prepared by [[pulse radiolysis]] techniques. It is mononuclear in dilute solution, and shows a strong absorption in the ultraviolet and a weaker absorption at 650 nm.<ref name="PulseRad"/> | ||
==References== | ==References== | ||
Revision as of 18:20, 26 June 2010
In inorganic nomenclature, a manganate is any negatively charged molecular 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).[2][3]
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.
- 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.[4][5]
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:[6]
- 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[7] with a visible absorption maximum of λmax = 670 nm (ε = 900 dm3 mol−1 cm−1).[2][3] 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).[7]
Hypomanganates may be prepared by the careful reduction of manganates with sulfite,[7] hydrogen peroxide[8] or mandelate.[3] 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:[6]
- HMnO2−4 ⇌ MnO3−4 + H+ pKa = 13.7 ± 0.2
Manganate(IV)
The manganate(IV) anion has been prepared by pulse radiolysis techniques. It is mononuclear in dilute solution, and shows a strong absorption in the ultraviolet and a weaker absorption at 650 nm.[6]
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.
- ↑ 2.0 2.1 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.
- ↑ 3.0 3.1 3.2 Lee, Donald G.; Chen, Tao Reduction of manganate(VI) by mandelic acid and its significance for development of a general mechanism of oxidation of organic compounds by high-valent transition metal oxides. J. Am. Chem. Soc. 1993, 115 (24), 11231–36. DOI: 10.1021/ja00077a023.
- ↑ 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.
- ↑ Sekula-Brzezińska, K.; Wrona, P. K.; Galus, Z. Rate of the MnO4−/MnO42− and MnO42−/MnO43− electrode reactions in alkaline solutions at solid electrodes. Electrochim. Acta 1979, 24 (5), 555–63. DOI: 10.1016/0013-4686(79)85032-X.
- ↑ 6.0 6.1 6.2 Rush, J. D.; Bielski, B. H. J. Studies of Manganate(V), -(VI), and -(VII) Tetraoxyanions by Pulse Radiolysis. Optical Spectra of Protonated Forms. Inorg. Chem. 1995, 34 (23), 5832–38. DOI: 10.1021/ic00127a022.
- ↑ 7.0 7.1 7.2 Greenwood, Norman N.; Earnshaw, A. Chemistry of the Elements; Pergamon: Oxford, 1984; pp 1221–22. ISBN 0-08-022057-6.
- ↑ Lee, Donald G.; Chen, Tao Oxidation of hydrocarbons. 18. Mechanism of the reaction between permanganate and carbon-carbon double bonds. J. Am. Chem. Soc. 1989, 111 (19), 7534–38. DOI: 10.1021/ja00201a039.
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