Solvated electron

A solvated electron is a free electron in (solvated in) a solution, and is the smallest possible anion. Solvated electrons occur widely, although it is difficult to observe them directly because their lifetimes are so short.[1] The deep color of solutions of alkali metals in liquid ammonia arises from the presence of solvated electrons: blue when dilute and copper-colored when more concentrated (> 3 molar).[2] Classically, 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 real hydration energy of the solvated electron can be estimated by using the hydration energy of a proton in water combined with kinetic data from pulse radiolysis experiments. The solvated electron forms an acid–base pair with atomic hydrogen.

The solvated electron is responsible for a great deal of radiation chemistry.

Alkali metals dissolve in liquid ammonia giving deep blue solutions, which conduct electricity. The blue colour of the solution is due to ammoniated electrons, which absorb energy in the visible region of light. Alkali metals also dissolve in some small primary amines, such as methylamine and ethylamine[3] and hexamethylphosphoramide, forming blue solutions.

Properties

Focusing on solutions in ammonia, liquid ammonia will dissolve all of the alkali metals and other electropositive metals such as Ca,[4] Sr, Ba, Eu, and Yb (also Mg using an electrolytic process[5]), giving characteristic blue solutions.

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-color phase becomes immiscible from a more dense blue phase. Above 8 MPM the solution is bronze/gold-colored. In the same concentration range the overall density decreases by 30%.

Dilute solutions are paramagnetic and at around 0.5 MPM all electrons are paired up and the solution becomes diamagnetic. Several models exist to describe the spin-paired species: as an ion trimer; as an ion-triple—a cluster of two single-electron solvated-electron species in association with a cation; or as a cluster of two solvated electrons and two solvated cations.

Solvated electrons produced by dissolution of reducing metals in ammonia and amines are the anions of salts called electrides. Such salts can be isolated by the addition of macrocyclic ligands such as crown ether and cryptands. These ligands bind strongly the cations and prevent their re-reduction by the electron.

Its standard electrode potential value is -2.77 V.[6] 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.[7]

Some thermodynamic properties of the solvated electron have been investigated by Joshua Jortner and Richard M. Noyes (1966)[8]

Alkaline aqueous solutions above pH = 9.6 regenerate the hydrated electron through the reaction of hydrated atomic hydrogen with hydroxide ion giving water beside hydrated electrons.

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.

The properties of solvated electron can be investigated using the rotating ring-disk electrode.

Reactivity and applications

The solvated electron reacts with oxygen to form a superoxide radical (O2.−).[9] With nitrous oxide, solvated electrons react to form hydroxyl radicals (HO.).[10] The solvated electrons can be scavenged from both aqueous and organic systems with nitrobenzene or sulfur hexafluoride.

A common use of sodium dissolved in liquid ammonia 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.

Solvated electrons are involved in the reaction of sodium metal with water.[11] Two solvated electrons combine to form molecular hydrogen and hydroxide ion.

Solvated electrons are also involved in electrode processes.[12]

Diffusion

The diffusivity of the solvated electron in liquid ammonia can be determined using potential-step chronoamperometry.[13]

In gas phase and upper atmosphere of Earth

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.[14]

Past

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). James Ballantyne Hannay and J. Hogarth repeated the experiments with sodium in 1879–1880. W. Weyl in 1844 and C. A. Seely in 1871 used liquid ammonia while Hamilton Cady in 1897 related the ionizing properties of ammonia to that of water. Charles A. Kraus measured the electrical conductance of metal ammonia solutions and in 1907 attributed it to the electrons liberated from the metal.[15][16] In 1918, G. E. Gibson and W. L. Argo introduced the solvated electron concept.[17] 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.[18]

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. Cotton, F. A.; Wilkinson, G. (1972). Advanced Inorganic Chemistry. John Wiley and Sons Inc. ISBN 978-0-471-17560-5.
  3. Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 978-0-08-037941-8.
  4. Edwin M. Kaiser (2001). "Calcium-Ammonia". Calcium–Ammonia. Encyclopedia of Reagents for Organic Synthesis. doi:10.1002/047084289X.rc003. ISBN 978-0471936237.
  5. 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.
  6. Baxendale, J. H. (1964), Radiation Res. Suppl., 114 and 139
  7. 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.
  8. 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.
  9. 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.
  10. 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.
  11. Walker, D.C. (1966). "Production of hydrated electron". Canadian Journal of Chemistry. 44 (18): 2226–. doi:10.1139/v66-336.
  12. B. E. Conway, D. J. MacKinnon, J. Phys. Chem., 74, 3663, 1970
  13. 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.
  14. F. Arnold, Nature 294, 732-733, (1981)
  15. 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.
  16. 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.
  17. 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.
  18. Dye, J. L. (2003). "Electrons as Anions". Science. 301 (5633): 607–608. doi:10.1126/science.1088103. PMID 12893933.

Further reading
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