Acidium, 132

In equimolar mixtures of nitric acid and water a monohydrate is formed whose Raman spectrum has been observed. There is no evidence for the existence of appreciable concentrations of the nitric acidium ion in aqueous nitric acid. 

The rates of nitration of mesitylene-a-sulphonate anion (iii) and iso-durene-a -sulphonate anion (iv) in mixtures of aqueous nitric and perchloric acid followed a zeroth-order rate law. Although the rate of exchange of oxygen could not be measured because of the presence of perchloric acid, these results again show that, under conditions most amenable to its existence and involvement, the nitric acidium ion is ineffective in nitration. 

If it be assumed that the ionising characteristics of nitric acid are similar to those of the organic indicators used to define the scales of acidity, then a correspondence between the acidity-dependence of nitration and would suggest the involvement of the nitronium ion, whereas a correspondence with Hq would support the h)rpothesis that the nitric acidium ion were active. The analogies with and Hg are expressed in the first and last pairs of the followii equations respectively. The symbol AQ represents anthraquinone, the indicator originally used in this way for comparison with the acidity dependence of the rate of nitration of nitrobenzene ... 

The nitric acidium ion undergoes slow heterolysis to yield water and the nitronium ion ... 

In the presence of sulphuric acid this route to the nitric acidium ion is almost entirely dominant, the additional route of protonation via the... 

Neither of the above schemes for forming the nitric acidium ion involves water. However, the addition of moderate quantities of water depresses the zeroth-order rate by up to a factor of four, without disturbing the kinetic form. This last fact shows that an inappreciable fraction of the nitronium ions is reacting with water, and therefore to explain the results it is necessary to postulate the existence of a means, involving water, for the consumption of nitric acidium ions ... 

The depletion of the concentration of nitric acidium ion by appreciable quantities of water is expressed by the equilibrium... 

NO3-] oc [N20J and so [NOai oc Now nitrate ions reduce the rate of formation of nitronium ion by de-protonating nitric acidium ions, and this effect must also depend upon [HN02]"toich> as was observed. 

In first-order nitration the anticatalysis is of the same form because the deprotonation of nitric acidium ion diminishes the stationary concentration of nitronium ion and therefore diminishes the rate of nitration. 

The anticatalytic action is ascribed to the deprotonation of nitric acidium ions by nitrite ions, which, being more basic than nitrate ions, will be more effective anticatalysts. The effect of nitrite ions should depend upon [HNOaJaioich it does. 

In summary, it is now more likely that the solvated nitrosyl ion, not the nitroso-acidium ion, is the nitrosating agent in diazotizations. 

Some interesting results have been obtained by Akand and Wyatt56 for the effect of added non-electrolytes upon the rates of nitration of benzenesulphonic acid and benzoic acid (as benzoic acidium ion in this medium) by nitric acid in sulphuric acid. Division of the rate coefficients obtained in the presence of nonelectrolyte by the concentration of benzenesulphonic acid gave rate coefficients which were, however, dependent upon the sulphonic acid concentration e.g. k2 was 0.183 at 0.075 molal, 0.078 at 0.25 molal and 0.166 at 0.75 molal (at 25 °C). With a constant concentration of non-electrolyte (sulphonic acid +, for example, 2, 4, 6-trinitrotoluene) the rate coefficients were then independent of the initial concentration of sulphonic acid and only dependent upon the total concentration of non-electrolyte. For nitration of benzoic acid a very much smaller effect was observed nitromethane and sulphuryl chloride had a similar effect upon the rate of nitration of benzenesulphonic acid. No explanation was offered for the phenomenon. 

The kinetic effect of increased pressure is also in agreement with the proposed mechanism. A pressure of 2000 atm increased the first-order rates of nitration of toluene in acetic acid at 20 °C and in nitromethane at 0 °C by a factor of about 2, and increased the rates of the zeroth-order nitrations of p-dichlorobenzene in nitromethane at 0 °C and of chlorobenzene and benzene in acetic acid at 0 °C by a factor of about 559. The products of the equilibrium (21a) have a smaller volume than the reactants and hence an increase in pressure speeds up the rate by increasing the formation of H2NO. Likewise, the heterolysis of the nitric acidium ion in equilibrium (22) and the reaction of the nitronium ion with the aromatic are processes both of which have a volume decrease, consequently the first-order reactions are also speeded up and to a greater extent than the zeroth-order reactions. 

The rate constants for the reaction remained relatively constant over the pH range 1-3.5, supporting the nitrous acidium ion mechanism. 

Secondary amines. The reaction of secondary amine type compounds with nitrous acid (HO NO) has been reviewed extensively by Turney and Wright (17), Ridd (18), Scanlan (19) and Mirvish (20). In a system containing (HO NO) as the nitrosating agent, the possible nitrosyl carriers are (H2O NO), (NO2 NO) and (NO+). The reactivity of (NO+) is very low and it is not considered an effective nitrosating form. Nitrous acidium ion (H2O NO+) plays a significant role only at concentrated acidic condi tions. Therefore, it seems likely that at the dilute acidic conditions that are encountered in the environment, it is nitrous anhydride (N2O3 = NO2 NO) which nitrosates secondary amines. 

The main nitrosating agent is probably nitrous acidium ion (NO OH2+). There is no simple rule, as there is for secondary amines, relating the ease of nitrosation to the properties of the amide. 

In der Essigsaure erscheinen als solvatisierteH -Ionen, die Acet-Acidium-ionen Hantzsch (74) ... 

An example may serve to illustrate the information that may be obtained from Fig. 10.1. Consider acetic acid. In water, acetic acid behaves as an acid or, to be more precise, an equimolar mixture of acetic acid and an acetate salt will have a pH of 4.74. If acetic acid is added to sulfuric acid, it will behave as a base and be leveled to CH3C(OH)j. the acetic acidium ion, and HSO4 (cf- Eq. 10.36 note the equilibrium lying at about -9 on the scale in Fig. 10.1). 

Note that an extension of the organic style of nomenclature as in (CH, CO.H ) - eltumoic acidium. is discouraged because it is based on the word acid and is often not easily adaptable to languages other than English.21... 

V (However, as long as " sulfuric acid" is Ihe only name used in the literature for IbSOj. "sulfuric acidium cation" (based an the previous IUPAC Red Book) will be necessary unless we use something trivial like "prolonalcd sulfuric acid. ]... 

Unfortunately, in spite of the importance of this kinetic form, the identity of the electrophile is not yet clearly established. The kinetic form (19) is consistent with nitrosation by either the nitrous acidium ion [eqn (20)] or the nitrosonium ion [eqn (21)]. 

The nitrosonium ion is a well-known chemical species and the equilibrium constant for its formation has been determined spectrophotometrically as 3 x 10-7 mol-1 dm3 (Bayliss et al., 1963).10 There is no spectroscopic evidence for the existence of the nitrous acidium ion. 

Thus the present evidence does not exclude the nitrosonium ion as the electrophile responsible for the kinetic form of equation (19) provided it is recognized that the half-life of the electrophile must then be extremely short. The nitrous acidium ion remains a possible electrophile but the complete absence of spectroscopic evidence for this ion weakens the case for its consideration (cf. Bayliss et al., 1963). 

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