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Ammonia in Solutions

The solubilities of ammonia in the lower aliphatic alcohols are considerable, although lower than those in water. Thus methyl alcohol dissolves 40 per cent, of ammonia at 0° C. The solubilities (under corresponding conditions of temperature and ammonia pressure) diminish with increasing molecular weight of the alcohol; in ethyl alcohol the solubilities are:

Temperature, ° C020
Solubility in grams of NH3 per litre of alcohol130.575
Density of solution0.7820.791

These solubilities are increased on the addition of water. Comparative solubilities of ammonia in ethyl, isopropyl, and isobutyl alcohols have also been determined at t° =20°±0.4° C.

Ammonia pressure in mm. of mercury457456.6523.1
Solubility, as volumes of ammonia gas dissolved by one volume of alcohol70.956.659.1

The following results show a diminution in the heat of solution of ammonia as the molar weight of the alcohol increases:

Range of concentration as percentage of NH30.17 to 9.940.30 to 4.23
Heat of solution in calories per mol. of NH3813 to 1508703 to 1230

The ratio in which ammonia is distributed between water and toluene is 26 at 19.5° C., and decreases with increasing concentration.

The distribution ratio between water and chloroform was redetermined by Moore and Winmill, who found 24.25 at 25° C. over the range of concentrations given in the table. The ratio alters considerably at higher concentrations according to Bell and Field. They find 24 in dilute solution, rising to 10 in very concentrated solution. Since this ratio is scarcely affected by the presence of ammonium chloride, the variation cannot be accounted for by a variation in the ionic equilibrium. It must be referred to the effect of increasing concentration on the relative proportions of the dissolved and hydrated ammonia molecules.

Since it is probable that ammonia dissolves in organic solvents of the water immiscible type as simple molecules, the distribution ratio between such solvents and water should give information as to the concentrations of simple molecules present in aqueous solution. This principle has been extensively used, e.g. to determine the hydrolytic equilibria of ammonium salts.

Conductivity and Dissociation Constant

The weakly alkaline reactions of solutions which contain dissolved ammonia are due to a slight electrolytic dissociation into NH4 and OH' ions. The approximate degrees of dissociation derived from the conductivities in the usual manner are:

Normality of solution0.010.101.0
Percentage dissociation410.3

The maximum conductivity is attained at a normality of about 3, at which k×104=11 (at 18° C.) and 13 (at 25° C.), and then falls to 1.93 in a 16 times normal solution (at 18° C.) and to 4.3 in a 13 times normal solution (at 25° C.). In the tables, the values of the conductivities at 18° C. are due to Kohlrausch, those at 25° C. to Goldschmidt, with the exception of the first two at the lowest concentrations, which are due to Bredig.

Conductivities of aqueous ammonia solutions (at 18° C.).

C in mols. per litre0.0590.2340.4670.9332.3074.558.8716.01

(at 25° C.)

C0.00390.00780.01090.02190.05530.11070.31480.541 0.6660.817

The mobility of the ammonium ion, L0, at zero concentration has been determined as follows:

Temperature, ° C182532.35

The molar conductivities at zero concentration, λ0, that is, the sums of the ionic conductivities of the ammonium and hydroxyl ion at zero (the lowest limiting) concentration, are as follows:

Temperature, °C1825

Ammonia is a weak or "half" electrolyte, the ionisation of which obeys the dilution law. Putting in the values of a=λ/λ0 derived from the conductivities into the equations

a constant is obtained which is found to hold good for the more dilute solutions. It must be carefully noted that the "undissociated ammonia," 1 - α, includes all forms not present as ions. The constant "k" is therefore an apparent one. It has the greatest practical utility, as it is available for calculating the alkalinities of ammonia solutions, as well as the acidities and degrees of hydrolysis of ammonium salts. The value of the constant of basic dissociation into ammonium and hydroxyl ions, kB at about 18° C., is 1.7 to 1.8×10-5. From the determination of λ and λ0 at other temperatures the following values of the constant are obtained:

Temperature, ° C.0182532.3550

At higher temperatures the constant diminishes; thus 105k is 1.35 at 100° C. and 0.0093 at 306° C.

The apparent constant also holds good, not only for free ammonia and its salts, but also for partly neutralised solutions of ammonia, so that from it the acidities "h" (or alkalinities) of any given mixture can be calculated by means of the equation

(kW is the water constant=1×10-14 at 23° C.).

The acidities of solutions of ammonium salts may be calculated from the equation

when kB=1.8×10-5.

At C. (normality)

These acidities correspond well with the turning-point of methyl red, and consequently this is a good indicator for the accurate titration of ammonia.

The converse calculation, that of the ionisation constants from the hydrolysis of ammonium salts, has been made by Lunden:

At C.018254060

Thermochemistry and Thermodynamics of Ammonia Solutions

The total (or integral) heat of solution of ammonia in water as determined by calorimetric methods is 8.430 Cals. for 1 mol. of ammonia in 200 mols. of water. The differential heat of solution (1 mol. in a large quantity of a solution having any given concentration) is also positive. It may be calculated by the Clapeyron equation from the ammonia pressures of, e.g., a dilute solution as above at 10° and 30° C., and is found to be 8.700 Cals. per mol. at 20°. This heat may be divided into (a) the latent heat of condensation of ammonia to the liquid form; (b) the heat of solution, hydration, ionisation.

(a) The latent heat of condensation to liquid is 5.000 Cals. in round numbers. From the vapour pressures, the latent heats at various temperatures have been calculated and expressed by the formula

This gives a molar latent heat of 4.890 Cals. at 15°. The differentiation of one formula of Cragoe, Meyers, and Taylor gives 3.950 Cals. Then approximately 8.430-5.000 = 3.430 Cals. are evolved in solution, hydration, and ionisation when liquid ammonia is dissolved in dilute solution. The heat of ionisation, Qi, is determined from the change in the dissociation constant with the temperature. Thus the mean value of Qi, between 18° and 25° is -1.190 Cals. (Moore and Winmill's results), while between 0° and 25° it is -1.675 (Kanolt's results). The heat of ionic dissociation is also obtained as the difference between the heats of neutralisation of (a) ammonia, Q1 and (b) a strong base, Q, with a strong acid; thus Qi=Q1-Q=12.300 – 13.800 = -1.500 Cals. (Q1 is 12.300 Cals. in the case of 0.278N ammonia, according to Thomsen.) From the results of distribution experiments between water and chloroform, and of conductivity determinations, the following heats may be deduced:

NH3+H2O = NH4OH + 5.450 Cals. (18° to 25° C.);
NH4OH = NH4+OH' - 3.630 Cals.

Hence, allowing for the proportions present in the solution, of NH3, NH3.xH2O and ammonium hydroxide (or any other kind of molecule which is in direct equilibrium with the ions), it is calculated that the heat of neutralisation in a 0.278N solution at 18° C. is 12.320 Cals., agreeing well with the value of Thomsen.

The Molecules present in Aqueous Ammonia

In the foregoing account of the solubility, dissociation, and thermochemistry of ammonia solutions, a judgment as to the kind of molecule which is produced by the hydration was left in suspense as far as possible. It is, however, convenient to use the term hydrate for that part of the ammonia which is in direct equilibrium with the gas, and the term hydroxide for that which is in equilibrium with the ions. The presence of one or more hydrates which may be more or less dissociated into NH3 and H2O is probable, and the presence of the ion NH4 in low concentration and perhaps also hydrated is demonstrated by electrochemical evidence. But whether an "ammonium hydroxide" analogous to sodium hydroxide exists in solution is open to doubt. It has been pointed out by Caven, that, according to the present electronic theory of valency, the ammonium ion is formed by the direct addition of the hydrion of water to NH3, and that the assumption of a non-ionised intermediate ammonium hydroxide is unnecessary. On the other hand, Moore has shown that the distribution ratios between water and chloroform, taken along with the electrolytic dissociation equilibria, require (that is, if the law of concentration action is assumed) a considerable proportion of ammonium hydroxide; so that the true dissociation constant k' is not the same as the apparent constant kB. The following table gives the values of the various molecular species calculated from these results. The constant kB is 1.725×10-5 (at 18° C.):

The molecules present in aqueous ammonia (Concentrations in mols. per c.c., ×105.)

Temp., °C.Total NH3 in any Form.NH3.xH2O. Hydrated Ammonia.NH4OH. Ammonium Hydroxide.NH4. Ammonium Ion.k2×105. Hydration Constant.k'B×105. Dissociation Constant.

The different views may be reconciled in various ways. The hydrion may attach itself to the NH4 which is already hydrated, giving the hydrated ion xH2O.NH4. On this hypothesis, NH3.xH2O in the above table may be substituted for NH4OH without affecting the quantitative relations; while for NH3.xH2O there would be substituted NH3 dissolved as such. It seems probable on the electronic theory that non-ionised NaOH, like non-ionised NaCl, does not exist, only the ions being present at all dilutions, because the electron has already passed from all the Na on combination. Only a portion of these ions, however, i.e. those which conduct the current, are dissociated. In the same sense we may suppose that non-ionised NH4OH does not exist; the NH3 which has attached a hydrogen ion is a positive ion at all dilutions. But a part of this is not dissociated, and takes the place of the older non-electrolyte, ammonium hydroxide, in calculations involving ammonium and hydroxyl-ion concentrations.

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