Editing Acid–base reaction (section)

===Solvent system definition===
One of the limitations of the Arrhenius definition is its reliance on water solutions. [[Edward Curtis Franklin]] studied the acid–base reactions in liquid ammonia in 1905 and pointed out the similarities to the water-based Arrhenius theory. [[Albert Fredrick Ottomar Germann|Albert F.O. Germann]], working with liquid [[phosgene]], {{chem|COCl|2}}, formulated the solvent-based theory in 1925, thereby generalizing the Arrhenius definition to cover aprotic solvents.<ref>{{cite journal |last=Germann |first=Albert F.O. |date=6 October 1925 |title=A General Theory of Solvent Systems |journal=[[Journal of the American Chemical Society]] |volume=47 |issue=10 |pages=2461–2468 |doi=10.1021/ja01687a006|bibcode=1925JAChS..47.2461G }}</ref>

Germann pointed out that in many solutions, there are ions in equilibrium with the neutral solvent molecules:
* [[onium ion|solvonium ions]]: a generic name for positive ions. These are also sometimes called solvo-acids;<ref name=SASB>{{Greenwood&Earnshaw1st|page=487}}</ref> when [[protonated]] solvent, they are [[lyonium ion]]s.
* [[ate complex|solvate ions]]: a generic name for negative ions. These are also sometimes called solve-bases;<ref name=SASB/> when deprotonated solvent, they are [[lyate ion]]s.

For example, water and [[ammonia]] undergo such dissociation into [[hydronium]] and [[hydroxide]], and [[ammonium]] and [[metal amides#Alkali metal amides|amide]], respectively:
<math chem display=block>\begin{align}
  \ce{2 H2O} & \ce{\, <=> H3O+ + OH-} \\[4pt]
  \ce{2 NH3} & \ce{\, <=> NH4+ + NH2-}
\end{align}</math>

Some aprotic systems also undergo such dissociation, such as [[dinitrogen tetroxide]] into [[nitrosonium]] and [[nitrate]],{{#tag:ref|Pure N<sub>2</sub>O<sub>4</sub> does not undergo such dissolution. However, it becomes electrically conductive when mixed with a polarized compound, which is believed to correspond with the establishment of such an equilibrium.<ref>{{Greenwood&Earnshaw1st|page=525}}</ref>|group=note}} [[antimony trichloride]] into dichloroantimonium and tetrachloroantimonate, and phosgene into chlorocarboxonium and [[chloride]]:
<math chem display=block>\begin{align}
  \ce{N2O4} & \ce{\, <=> NO+ + NO3-} \\[4pt]
  \ce{2 SbCl3} & \ce{\, <=> SbCl2+ + SbCl4-} \\[4pt]
  \ce{COCl2} & \ce{\, <=> COCl+ + Cl-}
\end{align}</math>

A solute that causes an increase in the concentration of the solvonium ions and a decrease in the concentration of solvate ions is defined as an ''acid''. A solute that causes an increase in the concentration of the solvate ions and a decrease in the concentration of the solvonium ions is defined as a ''base''.

Thus, in liquid ammonia, {{chem2|KNH2}} (supplying {{chem2|NH2-}}) is a strong base, and {{chem2|NH4NO3}} (supplying {{chem2|NH4+}}) is a strong acid. In liquid [[sulfur dioxide]] ({{chem2|SO2}}), [[thionyl]] compounds (supplying {{chem2|SO(2+)}}) behave as acids, and [[sulfites]] (supplying {{chem2|SO3(2-)}}) behave as bases.

The non-aqueous acid–base reactions in liquid ammonia are similar to the reactions in water:
<math chem display=block>\begin{align}
  \underset{\text{base}}{\ce{2 NaNH2}} + \underset{\text{amphiphilic} \atop \text{amide}}{\ce{Zn(NH2)2}} &\longrightarrow \ce{Na2[Zn(NH2)4]} \\[4pt]
  \underset{\text{acid}}{\ce{2 NH4I}} \ + \ \ce{Zn(NH2)2} &\longrightarrow \ce{[Zn(NH3)4]I2}
\end{align}</math>

Nitric acid can be a base in liquid sulfuric acid:
<math chem display=block>\underset{\text{base}}{\ce{HNO3}} + \ce{2 H2SO4 -> NO2+ + H3O+ + 2 HSO4-}</math>
The unique strength of this definition shows in describing the reactions in aprotic solvents; for example, in liquid {{chem2|N2O4}}:

<math chem display=block>\underset{\text{base}}{\ce{AgNO3}} + \underset{\text{acid}}{\ce{NOCl_{\ }}} \longrightarrow \underset{\text{solvent}}{\ce{N2O4}} + \underset{\text{salt}}{\ce{AgCl_{\ }}}</math>

Because the solvent system definition depends on the solute as well as on the solvent itself, a particular solute can be either an acid or a base depending on the choice of the solvent: {{chem2|HClO4}} is a strong acid in water, a weak acid in acetic acid, and a weak base in fluorosulfonic acid; this characteristic of the theory has been seen as both a strength and a weakness, because some substances (such as {{chem2|SO3}} and {{chem2|NH3}}) have been seen to be acidic or basic on their own right. On the other hand, solvent system theory has been criticized as being too general to be useful. Also, it has been thought that there is something intrinsically acidic about hydrogen compounds, a property not shared by non-hydrogenic solvonium salts.<ref name=review1940/>