Amphoteric Protic Solvents Undergo Solvent Autoionization just as Water Does The Brønsted-Lowry concept allows for an understanding of hydrogen ion transfer chemistry in amphoteric protic solvents. Amphoteric protic solvents are those which can both accept and receive hydrogen ions. From the viewpoint of the Brønsted-Lowry concept, the acid-base chemistry in these solvents is governed by autoionization equilibria analogous to water autoionization. \[\ce{2H_2O(l) ⇌ H_3O^+ + OH^{-}} \quad \quad K_w = 1.0\times10^{-14}\nonumber \] For example, sulfuric acid ionizes according to the equation: \[\ce{2H_2SO_4(l) ⇌ H_3SO_4^{+} + HSO_4^{-}} \quad \quad K = 4\times10^{-4}\nonumber \] The magnitude of the solvent autoionization constant in a given amphoteric solvent determines the amount of protonated and deprotonated* solvent present. Since sulfuric acid's autoionization constant is much larger than that of water (\(K_w = 10^{-14}\)), the concentrations of \(\ce{H_3SO_4^{+}}\) and \(\ce{HSO_4^{-}}\) present in pure sulfuric acid are ~0.02 M , much greater than the \(10^{-7}\) M \(H_3O^+\) and \(\ce{OH^{-}}\) present in pure water. In contrast, ammonia's autoionization constant is much less than that of water and only \(10^{-14}\, M\, \ce{NH_4^{+}}\) and \(\ce{NH_2^{-}}\) are present in pure ammonia. \[\ce{2NH_3(l) ⇌ NH_4^+ + NH_2^{-}} \quad \quad K = 10^{-27}\nonumber \] The Solvent Leveling Effect Limits the Strongest Acid or Base that Can Exist The conjugate acid and base of the solvent are the strongest Brønsted acids and bases that can exist in that solvent. To see why this is the case for acids, consider the reaction between a Brønsted acid (\(\ce{HA}\)) and solvent (\(\ce{S}\)): \[\ce{HA + S ⇌ HS^{+} + A^{-}}\nonumber \] This equilibrium will favor dissociation of whichever is a stronger acid - \(\ce{HA}\) or \(\ce{HS^{+}}\). If the acid is stronger it will mostly dissociate to give \(\ce{HS^{+}}\), while if the solvent's conjugate acid is stronger the acid will be mostly un-ionized and remain as \(\ce{HA}\).** Any acid significantly stronger than \(\ce{HS^{+}}\) will act as a strong acid and effectively dissociate completely to give the solvent's conjugate acid (\(\ce{HS^{+}}\)). This also means that the relative acidity of acids stronger than \(\ce{HS^{+}}\) cannot be distinguished in solvent \(\ce{S}\). This is called the leveling effect since the solvent "levels" the behavior of acids much stronger than itself to that of complete dissociation. For example, there is no way to distinguish the acidity of strong Arrhenius acids like \(\ce{HClO_4}\) and \(\ce{HCl}\) in water since they both completely dissociate. However, it is possible to distinguish their relative acidities in solvents that are more weakly basic than the conjugate base of the strongest\(†\) acid, since then the acids will dissociate to different extents. Such solvents are called differentiating solvents. For example, acetonitrile acts as a differentiating solvent towards \(\ce{HClO_4}\) and \(\ce{HCl}\). Both \(\ce{HClO_4}\) and \(\ce{HCl}\) partly dissociate in MeCN, with the stronger acid \(\ce{HClO_4}\) dissociating to a greater extent than \(\ce{HCl}\). The leveling effect can also occur in basic solutions. The strongest Brønsted base, \(\ce{B}\), that can exist in a solvent is determined by the relative acidity of the solvent, and the base's conjugate acid, \(\ce{BH^{+}}\), determines whether the base will remain unprotonated and able to act as a base in that solvent. If the solvent is represented this time as \(\ce{HS}\) then the relevant equilibrium is: \[\ce{B + HS ⇌ BH^{+} + S^{-}} \nonumber \] The position of this equilibrium depends on whether \(\ce{B}\) or \(\ce{S^{-}}\) is the stronger base. If the solvent's conjugate base, \(\ce{S^{-}}\), is stronger, then the base B will remain unprotonated and available to act as a base. However, if \(\ce{B}\) is a stronger base than \(\ce{S^{-}}\), it will deprotonate the solvent to give \(\ce{BH^{+}}\) and \(\ce{S^{-}}\). In this way the strongest base that can exist in a given solvent is the solvent's conjugate base. The relative strength of Brønsted bases can only be determined in solvents that are more weakly acidic than \(\ce{BH^{+}}\); otherwise the bases will all be leveled to \(\ce{S^{-}}\). It is important to consider the leveling effect of protic solvents when performing syntheses that require the use of basic reagents. For instance, hydride and carbanion reagents (lithium aluminum hydride, Grignard reagents, alklyllithium reagents, etc.) cannot be used as nucleophiles in protic solvents like water, alcohols, or enolizable aldehydes and ketones. Since carbanions are stronger bases than these solvents' conjugate bases, they will instead act as Brønsted bases and deprotonate the solvent. For example if one adds n-butyllithium to water, the result (along with much heat and possibly a fire) will be butane and a solution of lithium hydroxide: \[\ce{Li^{+} + ^{-}:CH_2CH_2CH_2CH_3 + H_2O(l) → Li^{+}(aq) + OH^{-}(aq) + CH_3CH_2CH_2CH_3(g)} \nonumber \] Since hydride and carbanion reagents cannot be used as nucleophiles in protic solvents like water or methanol, they are commonly sold as solutions in solvents such as hexanes (for alklyllithium reagents) or tetrahydrofuran (for Grignard reagents and \(\ce{LiAlH_4}\)).