Solvated electron

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A solvated electron is a free electron in (solvated in) a solution, and is the smallest possible anion. Solvated electrons occur widely although they are often not observed directly. The deep color of solutions of alkali metals in ammonia arises from the presence of solvated electrons: blue when dilute and copper-colored when more concentrated (> 3 molar).[1] Classically, discussions of solvated electrons focus on their solutions in ammonia, which are stable for days, but solvated electrons occur in water and other solvents, in fact, any solvent that mediates outer-sphere electron transfer. The solvated electron is responsible for a great deal of radiation chemistry.

Properties[edit]

Focusing on ammonia solutions, all of the alkali metals, as well as Ca, Sr, Ba, Eu, and Yb, dissolve to give the characteristic blue solutions. Other amines, such as methylamine and ethylamine, are also suitable solvents.[2]

A lithium ammonia solution at −60 °C is saturated at about 16 mol% metal (16 MPM in the local jargon). 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, or 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.

Reactivity and applications[edit]

The solvated electron reacts with oxygen to form a superoxide radical (O2.-),[3] which is a potent oxidant. With nitrous oxide, solvated electrons react to form hydroxyl radicals (HO.).[4] The solvated electrons can be scavenged from both aqueous and organic systems with nitrobenzene or sulfur hexafluoride[citation needed].

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.

History[edit]

The first to note the color of metal-electride solutions was Sir 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 were the first to use liquid ammonia. Hamilton Cady in 1897 was the first to relate the ionizing properties of ammonia to that of water. Charles A. Kraus measured the electrical conductance of metal ammonia solutions and in 1907 was the first to attribute it the electrons liberated from the metal.[5][6] In 1918, G. E. Gibson and W. L. Argo introduced the solvated electron concept.[7] 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.[8]

References[edit]

  1. ^ Cotton, F.A; G. Wilkinson (1972). Advanced Inorganic Chemistry. John Wiley and Sons Inc. ISBN 0-471-17560-9. 
  2. ^ Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 0080379419. 
  3. ^ Susan E. Forest, Michael J. Stimson, and John D. Simon. "Mechanism for the Photochemical Production of Superoxide by Quinacrine". J. Phys. Chem. B1999,103,3963-3964
  4. ^ Janata, E.; Schuler, Robert H. (1982). "Rate constant for scavenging eaq- in nitrous oxide-saturated solutions". J. Phys. Chem. 86 (11): 2078–84. 
  5. ^ Solutions of Metals in Non-Metallic Solvents; I. General Properties of Solutions of Metals in Liquid Ammonia Charles A. Kraus J. Am. Chem. Soc., 1907, 29 (11), pp 1557–1571 doi:10.1021/ja01965a003
  6. ^ A Molecular Perspective on Lithium–Ammonia Solutions Eva Zurek, Peter P. Edwards, and Roald Hoffmann Angew. Chem. Int. Ed. 2009, vol. 48, 8198 – 8232 doi:10.1002/anie.200900373
  7. ^ The Absorption Spectra of the Blue Solutions of Certain Alkali and Alkaline Earth Metals in Liquid Ammonia and Methylamine G. E. Gibson, W. L. Argo J. Am. Chem. Soc., 1918, vol. 40, pp 1327–1361. doi:10.1021/ja02242a003
  8. ^ Dye, J. L. (2003). "Electrons as Anions". Science 301 (5633): 607–608. doi:10.1126/science.1088103. PMID 12893933. 

9) Hydrated Electrons at the Water/Air Interface; DOI: 10.1021/ja101176r

             D. M. Sagar, Colin. D. Bain and Jan R. R. Verlet,J. Am. Chem. Soc., 2010, 132 (20), pp 6917–6919.