Chemistry:Solvated electron

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Short description: Free electron in a solution, often liquid ammonia

A solvated electron is a free electron in (solvated in) a solution, and is the smallest possible anion. Solvated electrons occur widely.[1] Often, discussions of solvated electrons focus on their solutions in ammonia, which are stable for days, but solvated electrons also occur in water and other solvents – in fact, in any solvent that mediates outer-sphere electron transfer. The solvated electron is responsible for a great deal of radiation chemistry.

Ammonia solutions

Liquid ammonia will dissolve all of the alkali metals and other electropositive metals such as Ca,[2] Sr, Ba, Eu, and Yb (also Mg using an electrolytic process[3]), giving characteristic blue solutions. For alkali metals in liquid ammonia, the solution is blue when dilute and copper-colored when more concentrated (> 3 molar).[4] These solutions conduct electricity. The blue colour of the solution is due to ammoniated electrons, which absorb energy in the visible region of light. The diffusivity of the solvated electron in liquid ammonia can be determined using potential-step chronoamperometry.[5]

Solvated electrons in ammonia are the anions of salts called electrides.

Na + 6 NH3 → [Na(NH3)6]+e

The reaction is reversible: evaporation of the ammonia solution produces a film of metallic sodium.

Case study: Li in NH3

Photos of two solutions in round-bottom flasks surrounded by dry ice; one solution is dark blue, the other golden.
Solutions obtained by dissolution of lithium in liquid ammonia. The solution at the top has a dark blue color and the lower one a golden color. The colors are characteristic of solvated electrons at electronically insulating and metallic concentrations, respectively.

A lithium–ammonia solution at −60 °C is saturated at about 15 mol% metal (MPM). When the concentration is increased in this range electrical conductivity increases from 10−2 to 104 ohm−1cm−1 (larger than liquid mercury). At around 8 MPM, a "transition to the metallic state" (TMS) takes place (also called a "metal-to-nonmetal transition" (MNMT)). At 4 MPM a liquid-liquid phase separation takes place: the less dense gold-colored phase becomes immiscible from a denser blue phase. Above 8 MPM the solution is bronze/gold-colored. In the same concentration range the overall density decreases by 30%.

Other solvents

Alkali metals also dissolve in some small primary amines, such as methylamine and ethylamine[6] and hexamethylphosphoramide, forming blue solutions. THF dissolves alkali metal, but a Birch reduction (see § Applications) analogue does not proceed without a diamine ligand.[7] Solvated electron solutions of the alkaline earth metals magnesium, calcium, strontium and barium in ethylenediamine have been used to intercalate graphite with these metals.[8]

Water

Solvated electrons are involved in the reaction of alkali metals with water, even though the solvated electron has only a fleeting existence.[9] Below pH = 9.6 the hydrated electron reacts with the hydronium ion giving atomic hydrogen, which in turn can react with the hydrated electron giving hydroxide ion and usual molecular hydrogen H2.[10]

Solvated electrons can be found even in the gas phase. This implies their possible existence in the upper atmosphere of Earth and involvement in nucleation and aerosol formation.[11]

Its standard electrode potential value is -2.77 V.[12] Equivalent conductivity 177 Mho cm2 is similar to that of hydroxide ion. This value of equivalent conductivity corresponds to a diffusivity of 4,75*10−5 cm2s−1.[13]

Reactivity

Although quite stable, the blue ammonia solutions containing solvated electrons degrade rapidly in the presence of catalysts to give colorless solutions of sodium amide:

2 [Na(NH3)6]+e → H2 + 2 NaNH2 + 10 NH3

Electride salts can be isolated by the addition of macrocyclic ligands such as crown ether and cryptands to solutions containing solvated electrons. These ligands strongly bind the cations and prevent their re-reduction by the electron.

[Na(NH3)6]+e + cryptand → [Na(cryptand)]+e+ 6 NH3

The solvated electron reacts with oxygen to form a superoxide radical (O2.−).[14] With nitrous oxide, solvated electrons react to form hydroxyl radicals (HO.).[15]

Applications

Solvated electrons are involved in electrode processes, a broad area with many technical applications (electrosynthesis, electroplating, electrowinning).

A specialized use of sodium-ammonia solutions is the Birch reduction. Other reactions where sodium is used as a reducing agent also are assumed to involve solvated electrons, e.g. the use of sodium in ethanol as in the Bouveault–Blanc reduction.

History

The observation of the color of metal-electride solutions is generally attributed to Humphry Davy. In 1807–1809, he examined the addition of grains of potassium to gaseous ammonia (liquefaction of ammonia was invented in 1823).[16] James Ballantyne Hannay and J. Hogarth repeated the experiments with sodium in 1879–1880.[17] W. Weyl in 1864 and C. A. Seely in 1871 used liquid ammonia, whereas Hamilton Cady in 1897 related the ionizing properties of ammonia to that of water.[18][19][20] Charles A. Kraus measured the electrical conductance of metal ammonia solutions and in 1907 attributed it to the electrons liberated from the metal.[21][22] In 1918, G. E. Gibson and W. L. Argo introduced the solvated electron concept.[23] They noted based on absorption spectra that different metals and different solvents (methylamine, ethylamine) produce the same blue color, attributed to a common species, the solvated electron. In the 1970s, solid salts containing electrons as the anion were characterized.[24]

References

  1. Schindewolf, U. (1968). "Formation and Properties of Solvated Electrons". Angewandte Chemie International Edition in English 7 (3): 190–203. doi:10.1002/anie.196801901. 
  2. Edwin M. Kaiser (2001). "Calcium–Ammonia". Encyclopedia of Reagents for Organic Synthesis. doi:10.1002/047084289X.rc003. ISBN 978-0471936237. 
  3. Combellas, C; Kanoufi, F; Thiébault, A (2001). "Solutions of solvated electrons in liquid ammonia". Journal of Electroanalytical Chemistry 499: 144–151. doi:10.1016/S0022-0728(00)00504-0. 
  4. Cotton, F. A.; Wilkinson, G. (1972). Advanced Inorganic Chemistry. John Wiley and Sons Inc. ISBN 978-0-471-17560-5. 
  5. Harima, Yutaka; Aoyagui, Shigeru (1980). "The diffusion coefficient of solvated electrons in liquid ammonia". Journal of Electroanalytical Chemistry and Interfacial Electrochemistry 109 (1–3): 167–177. doi:10.1016/S0022-0728(80)80115-X. 
  6. Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 978-0-08-037941-8. 
  7. Burrows, James; Kamo, Shogo; Koide, Kazunori (2021-11-05). "Scalable Birch reduction with lithium and ethylenediamine in tetrahydrofuran". Science 374 (6568): 741–746. doi:10.1126/science.abk3099. ISSN 0036-8075. https://doi.org/10.1126/science.abk3099. 
  8. W. Xu and M. M. Lerner, "A New and Facile Route Using Electride Solutions To Intercalate Alkaline Earth Ions into Graphite", Chem. Mater. 2018, 30, 19, 6930–6935. https://doi.org/10.1021/acs.chemmater.8b03421
  9. Walker, D.C. (1966). "Production of hydrated electron". Canadian Journal of Chemistry 44 (18): 2226–. doi:10.1139/v66-336. 
  10. Jortner, Joshua; Noyes, Richard M. (1966). "Some Thermodynamic Properties of the Hydrated Electron". The Journal of Physical Chemistry 70 (3): 770–774. doi:10.1021/j100875a026. 
  11. F. Arnold, Nature 294, 732-733, (1981)
  12. Baxendale, J. H. (1964), Radiation Res. Suppl., 114 and 139
  13. Hart, Edwin J. (1969). "The Hydrated Electron". Survey of Progress in Chemistry 5: 129–184. doi:10.1016/B978-0-12-395706-1.50010-8. ISBN 9780123957061. 
  14. Hayyan, Maan; Hashim, Mohd Ali; Alnashef, Inas M. (2016). "Superoxide Ion: Generation and Chemical Implications". Chemical Reviews 116 (5): 3029–3085. doi:10.1021/acs.chemrev.5b00407. PMID 26875845. 
  15. Janata, Eberhard; Schuler, Robert H. (1982). "Rate constant for scavenging eaq- in nitrous oxide-saturated solutions". The Journal of Physical Chemistry 86 (11): 2078–2084. doi:10.1021/j100208a035. 
  16. Thomas, Sir John Meurig; Edwards, Peter; Kuznetsov, Vladimir L. (January 2008). "Sir Humphry Davy: Boundless Chemist, Physicist, Poet and Man of Action". ChemPhysChem 9 (1): 59–66. "An entry from Humphry Davy′s laboratory notebook of November 1808. It reads “When 8 Grains of potassium were heated in ammoniacal gas—it assumed a beautiful metallic appearance & gradually became of a fine blue colour”.". 
  17. Hannay, J. B.; Hogarth, James (26 February 1880). "On the solubility of solids in gases". Proceedings of the Royal Society of London 30 (201): 178–188. https://www.biodiversitylibrary.org/item/139575#page/202/mode/1up. 
  18. Weyl, W. (1864). "Ueber Metallammonium-Verbindungen" (in German). Annalen der Physik und Chemie 121: 601–612. https://babel.hathitrust.org/cgi/pt?id=mdp.39015065833884&view=1up&seq=621&skin=2021. 
  19. Seely, Charles A. (14 April 1871). "On ammonium and the solubility of metals without chemical action". The Chemical News 23 (594): 169–170. https://babel.hathitrust.org/cgi/pt?id=uc1.$c193335&view=1up&seq=177&skin=2021. 
  20. Cady, Hamilton P. (1897). "The electrolysis and electrolytic conductivity of certain substances dissolved in liquid ammonia". The Journal of Physical Chemistry 1: 707–713. https://babel.hathitrust.org/cgi/pt?id=mdp.39015026388507&view=1up&seq=737&skin=2021. 
  21. Kraus, Charles A. (1907). "Solutions of metals in non-metallic solvents; I. General properties of solutions of metals in liquid ammonia". J. Am. Chem. Soc. 29 (11): 1557–1571. doi:10.1021/ja01965a003. https://zenodo.org/record/1428868. 
  22. Zurek, Eva (2009). "A molecular perspective on lithium–ammonia solutions". Angew. Chem. Int. Ed. 48 (44): 8198–8232. doi:10.1002/anie.200900373. PMID 19821473. 
  23. Gibson, G. E.; Argo, W. L. (1918). "The absorption spectra of the blue solutions of certain alkali and alkaline earth metals in liquid ammonia and methylamine". J. Am. Chem. Soc. 40 (9): 1327–1361. doi:10.1021/ja02242a003. https://zenodo.org/record/1429048. 
  24. Dye, J. L. (2003). "Electrons as anions". Science 301 (5633): 607–608. doi:10.1126/science.1088103. PMID 12893933. 

Further reading