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 it directly since their lifetimes are so short. 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). 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 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 are conducting in nature. 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 hexamethylphosphoramide, forming blue solutions.
Focusing on ammonia solutions, all of the alkali metals, as well as Ca, Sr, Ba, Eu, and Yb (also Mg using an electrolytic process), dissolve to give the characteristic blue solutions. Other amines, such as methylamine and ethylamine, are also suitable solvents.
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. 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.
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 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 applicationsEdit
The solvated electron reacts with oxygen to form a superoxide radical (O2.−). With nitrous oxide, solvated electrons react to form hydroxyl radicals (HO.). 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.
The first recorded observation of the color of metal-electride solutions is generally attributed to 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. In 1918, G. E. Gibson and W. L. Argo introduced the solvated electron concept. 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.
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