Nitric acid, dissociation

The protonated nitric acid dissociates to form a nitronium ion (+N02). 

The constants k and k are independent of the medium and of the concentration of sulphuric acid. If Q denotes the proportion of nitric acid dissociated into N02+ ions, then ... 

This simplification frees us from the need to estimate the aqueous-phase concentration Caq (z, t). This will be a rather good approximation for a very soluble species, that is, a species with sufficiently high effective Henry s law constant H. Nitric acid, with a Henry s law constant of 2.1 x 105 M atm-1, is a good example of such a species. Recalling that dissolved nitric acid dissociates to produce nitrate, we obtain... 

This leaves the second order rate coefficient for toluene nitration in aqueous solution (k) as the only missing parameter. The most reliable data are probably those of Coombes et al(25). Unfortunately, they used nitric acid concentrations much below those found in heterogeneous nitrations and there is some uncertainty as to how they should best be extrapolated. Thus Cox and Strachan(9 ) have attributed deviations in their results to the probability of k values in their nitrating mixtures being higher than those measured by Coombes et al. Schiefferle(26) found it necessary to modify the acidity function concept to account for the contribution of the nitric acid dissociation to the proton concentration and suggested that k should be related to the total acidity rather than just the sulphuric acid concentration. A model based on this concept gave a better fit to his experimental results. 

HNOaCg) HN03(aq) the nitric acid dissociates readily to nitrate,... 

Therefore, in the cases of both additives, the kinetic law for the catalysis will assume a linear form when the concentration of the added species, or, in the case of sulphuric acid, the nitronium ion generated by its action, is comparable with the concentration of the species already present. This effect was observed to occur when the concentration of additive was about o-2 mol 1, a value in fair agreement with the estimated degree of dissociation of nitric acid ( 2.2.1). 

Acid—Base Chemistry. Acetic acid dissociates in water, pK = 4.76 at 25°C. It is a mild acid which can be used for analysis of bases too weak to detect in water (26). It readily neutralizes the ordinary hydroxides of the alkaU metals and the alkaline earths to form the corresponding acetates. When the cmde material pyroligneous acid is neutralized with limestone or magnesia the commercial acetate of lime or acetate of magnesia is obtained (7). Acetic acid accepts protons only from the strongest acids such as nitric acid and sulfuric acid. Other acids exhibit very powerful, superacid properties in acetic acid solutions and are thus useful catalysts for esterifications of olefins and alcohols (27). Nitrations conducted in acetic acid solvent are effected because of the formation of the nitronium ion, NO Hexamethylenetetramine [100-97-0] may be nitrated in acetic acid solvent to yield the explosive cycl o trim ethyl en etrin itram in e [121 -82-4] also known as cyclonit or RDX. 

Hexafluorophosphoric Acid. Hexafluorophosphoric acid (3) is present under ambient conditions only as an aqueous solution because the anhydrous acid dissociates rapidly to HF and PF at 25°C (56). The commercially available HPF is approximately 60% HPF based on PF analysis with HF, HPO2F2, HPO F, and H PO ia equiUbrium equivalent to about 11% additional HPF. The acid is a colorless Hquid which fumes considerably owiag to formation of an HF aerosol. Frequently, the commercially available acid has a dark honey color which is thought to be reduced phosphate species. This color can be removed by oxidation with a small amount of nitric acid. When the hexafluorophosphoric acid is diluted, it slowly hydrolyzes to the other fluorophosphoric acids and finally phosphoric acid. In concentrated solutions, the hexafluorophosphoric acid estabUshes equiUbrium with its hydrolysis products ia relatively low concentration. Hexafluorophosphoric acid hexahydrate [40209-76-5] 6 P 31.5°C, also forms (66). This... 

Peroxonitrous acid can decompose by two pathways isomerization to nitric acid, and dissociation into the hydroxyl radical and nitrogen dioxide. 

Aluminum nitrate is available commercially as aluminum nitrate nonahydrate [7784-27-2], A1(N02)3 9H20. It is a white, crystalline material with a melting point of 73.5°C that is soluble in cold water, alcohols, and acetone. Decomposition to nitric acid [7699-37-2], HNO, and basic aluminum nitrates [13473-90-0], A1(0H) (N03) where x + = 3, begins at 130°C, and dissociation to aluminum oxide and oxides of nitrogen occurs above 500°C. 

Amino-2-hydroxybenZOiC acid. This derivative (18) more commonly known as 4-aminosa1icy1ic acid, forms white crystals from ethanol, melts with effervescence and darkens on exposure to light and air. A reddish-brown crystalline powder is obtained on recrystallization from ethanol —diethyl ether. The compound is soluble ia dilute solutioas of nitric acid and sodium hydroxide, ethanol, and acetone slightly soluble in water and diethyl ether and virtually insoluble in benzene, chloroform or carbon tetrachloride. It is unstable in aqueous solution and decarboxylates to form 3-amiaophenol. Because of the instabihty of the free acid, it is usually prepared as the hydrochloride salt, mp 224 °C (dec), dissociation constant p 3.25. 

Qua.driva.Ient, Zirconium tetrafluoride is prepared by fluorination of zirconium metal, but this is hampered by the low volatility of the tetrafluoride which coats the surface of the metal. An effective method is the halogen exchange between flowing hydrogen fluoride gas and zirconium tetrachloride at 300°C. Large volumes are produced by the addition of concentrated hydrofluoric acid to a concentrated nitric acid solution of zirconium zirconium tetrafluoride monohydrate [14956-11-3] precipitates (69). The recovered crystals ate dried and treated with hydrogen fluoride gas at 450°C in a fluid-bed reactor. The thermal dissociation of fluorozirconates also yields zirconium tetrafluoride. 

Arsenic Acids and the Arsenates. Commercial arsenic acid, corresponds to the composition, one mole of arsenic pentoxide to four moles of water, and probably is the arsenic acid hemihydrate [7774-41-6] H AsO O.5H2O. It is obtained by treatment of arsenic trioxide with concentrated nitric acid. Solutions of this substance or of arsenic pentoxide in water behave as triprotic acids with successive dissociation constants = 5.6 x 10 , ... 

The existence of the nitronium ion in sulfuric-nitric acid mixtures was demonstrated both by cryoscopic measurements and by spectroscopy. An increase in the strong acid concentration increases the rate of reaction by shifting the equilibrium of step 1 to the right. Addition of a nitrate salt has the opposite effect by suppressing the preequilibrium dissociation of nitric acid. It is possible to prepare crystalline salts of nitronium ions, such as nitronium tetrafluoroborate. Solutions of these salts in organic solvents rapidly nitrate aromatic compounds. ... 

Examples include hydrochloric acid, nitric acid, and sulphuric acid. These are strong acids which are almost completely dissociated in water. Weak acids, such as hydrogen sulphide, are poorly dissociated producing low concentrations of hydrogen ions. Acids tend to be coiTosive with a sharp, sour taste and turn litmus paper red they give distinctive colour changes with other indicators. Acids dissolve metals such as copper and liberate hydrogen gas. They also react with carbonates to liberate carbon dioxide ... 

The Dissociation Constant of Nitric Acid. The largest value of K in Table 9 is that for the (HS04) ion. In Fig. 36 there is a gap of more than 0.2 electron-volt below the level of the (H30)1 ion. As is well known, several acids exist which in aqueous solution fall iu the intermediate region between the very weak acids and the recognized strong acids the proton levels of these acids will fall in this gap. The values of K for these acids obtained by different methods seldom show close agreement. Results obtained by various methods were compared in 1946 by Redlich,1 who discussed the difficulties encountered. 

At higher concentrations the Raman spectra of aqueous solutions of alkali nitrates and of nitric acid have been investigated. Nitric acid was found to be incompletely dissociated, though for the alkali nitrates no evidence of incomplete dissociation was found. Since accurate measurements on solutions of nitric acid have not been made at concentrations below 4.0 molar, it is not certain how the extrapolation to infinite... 

Using (204), let us attempt to predict the degree of dissociation of nitric acid in methanol solution. According to (204) the occupied proton... 

The ionisation may be attributed to the great tendency of the free hydrogen ions H+ to combine with water molecules to form hydroxonium ions. Hydrochloric and nitric acids are almost completely dissociated in aqueous solution in accordance with the above equations this is readily demonstrated by freezing-point measurements and by other methods. 

If the dissociation constant of the acid HA is very small, the anion A- will be removed from the solution to form the undissociated acid HA. Consequently more of the salt will pass into solution to replace the anions removed in this way, and this process will continue until equilibrium is established (i.e. until [M + ] x [A-] has become equal to the solubility product of MA) or, if sufficient hydrochloric acid is present, until the sparingly soluble salt has dissolved completely. Similar reasoning may be applied to salts of acids, such as phosphoric(V) acid (K1 = 7.5 x 10-3 mol L-1 K2 = 6.2 x 10-8 mol L-1 K3 = 5 x 10 13 mol L-1), oxalic acid (Kx = 5.9 x 10-2 mol L-K2 = 6.4 x 10-5molL-1), and arsenic)V) acid. Thus the solubility of, say, silver phosphate)V) in dilute nitric acid is due to the removal of the PO ion as... 

The amount of reddish-purple acid-chloranilate ion liberated is proportional to the chloride ion concentration. Methyl cellosolve (2-methoxyethanol) is added to lower the solubility of mercury(II) chloranilate and to suppress the dissociation of the mercury(II) chloride nitric acid is added (concentration 0.05M) to give the maximum absorption. Measurements are made at 530nm in the visible or 305 nm in the ultraviolet region. Bromide, iodide, iodate, thiocyanate, fluoride, and phosphate interfere, but sulphate, acetate, oxalate, and citrate have little effect at the 25 mg L 1 level. The limit of detection is 0.2 mg L 1 of chloride ion the upper limit is about 120 mg L . Most cations, but not ammonium ion, interfere and must be removed. 

The Raman spectrum of the monohydrate, HN03.H20, shows it to exist as the hydroxoni-um salt, H30+N03 13. Also, according to analyses of the Raman spectrum, nitric acid exists in aq solns either as a pseudo-acid, N02.0H or as a true acid, N03".H+. In 10 molar aq soln, both acids are present in equal amounts, being caused by the self-dissociation of nitrogen pentoxide (NjOj), while in a 6 molar soln, the pseudo acid is present only to the extent of 2%. and the more dilute the soln, the less pseudo acid is present. In very coned solns, the true acid is present only in small quantities (Refs 32 33)... 

The reflux of aqueous Pu(IV) solutions containing <6 M HNO3 produces polymer precipitates that are resistant to subsequent dissociation and dissolution in nitric acid. Eapid aging of the Pu(IV) polymer to form a PuC -like structure is responsible for the unusually stable polymer. Comparative studies under nonreflux conditions show that polymer does not form at concentrations of HNO3 >3 M. 

Concentrated nitric acid can effect nitration but it is not as reactive as a mixture of nitric acid with sulfuric acid. The active nitrating species in both media is the nitronium ion, NOz+, which is formed by protonation and dissociation of nitric acid. The concentration of NOz+ is higher in the more strongly acidic sulfuric acid than in nitric acid. 

It can be seen from Figure 5.18 that the KD values for zirconium are higher than those for hafnium at all nitric acid concentrations. This is because the dissolution of zirconium nitrate (Zr(N03)4) into zirconyl (Zr02+) and nitrate (NOj) ions takes place to a lower extent as compared to the corresponding dissolution of hafnium nitrate in an aqueous medium. Hence, separation is feasible. However, at higher nitric acid concentrations the separation factor is reduced significantly because the dissociation of hafnium nitrate (Hf(NOs)4) decreases sharply with increasing nitric acid concentration, with the result that the separation factor, p, falls off rapidly. Hence, the separation process calls for the adjustment of the nitric acid concentration to a suitably low value. 

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