Liquid Ammonia

Ammonia was first liquefied by Faraday, who heated a compound of silver chloride and ammonia contained in one limb of a bent tube, the other limb being cooled in ice. Bunsen showed that it could be liquefied under atmospheric pressure in a good freezing mixture (ice and calcium chloride), and hence gave an estimate of its boiling-point, which has since been determined by many investigators.

Boiling-point of anhydrous ammonia ° C at normal pressure: from -33.46 to -33.2.

The melting-point of solid ammonia has been variously given as:

-78.2°, -77.7°, and -77.0° C.

The specific heat of liquid ammonia has been determined with the following results:

Specific heat of liquid ammonia

Temperature Interval, ° C.	Specific Heat.
0 to +40	0.88 to 0.89
30 to +62	1.22876
0 to +30	1.0206
-188 to -103	0.50
-40	1.054
0	1.099
+40	1.162
0 to +20	1.152
+20 to +50	1.172
+30	1.168
50	1.222
70	1.297
90	1.131
100	1.538

According to Dieterici, the specific heats between 0° and 70° C. are:

t°C.	0	10	20	30	40	50	60	70
Spec. heat.	1.118	1.140	1.161	1.181	1.203	1.225	1.247	1.269

Densities of liquid ammonia and its saturated vapour

Temperature, °C.	-50	-40	-30	-20	-10	0	+10	+20	+40	+60	+80	+100
Density	0.6954	0.6835	0.6715	0.6593	0.6469	0.6341	0.6207	0.6067	0.5756	0.5404	0.5004	0.4522

The isobaric volumes and densities of liquid and vapour under a pressure equal to that of the vapour have been determined by Dieterici and Berthoud.

Isobaric specific volumes of liquid ammonia and its vapour

Temperature, °C.	0	+10	+20	+30	+40	+50	+60	+70	+80	+90	+100
vL	1.5656	1.5985	1.6342	1.6765	1.7227	1.7719	1.8250	1.8875	1.9595	2.0390	2.1525
vV	294.0	206.0	148.0	108.5	80.9	62.3	48.6	37.8	30.2	24.4	17.05

Isobaric densities of liquid ammonia and its vapour

Temperature, ° C.	0	+45.0	78.7	98.75	109.2	116.4	121.3	123.2	125.4	129.6
DL	0.6389	0.5696	0.5120	0.4640	0.4339	0.4056	0.3831	0.3750	0.3584	0.3246
DV	0.0034	0.0150	0.0322	0.0533	0.0691	0.0873	0.1024	0.1085	0.1220	0.1509

The vapour pressures of liquid ammonia were first determined by Faraday, and afterwards by Bunsen, Pictet, and others, but their results are now merely of historical interest.

These constants have been redetermined with great accuracy by methods which deserve mention, as they are widely used in similar cases. The apparatus is described in a paper by Cragoe, Meyers, and Taylor. A static method was employed. The pressures were read on mercury manometers, the whole being immersed in thermostats containing gasolene which could be regulated to 0.01° C.; or better, at the lower temperatures by means of coils in which carbon dioxide was allowed to expand. It was shown that the presence of incondensable gases caused lag in the attainment of the equilibrium pressures. Therefore the greatest care was taken to purify the ammonia. A large quantity of synthetic ammonia was distilled twice, the middle portion of the second distillate was allowed to stand in contact with metallic sodium for about a week, the liquid was shaken and hydrogen blown off. After several distillations in a high-pressure apparatus, the middle fraction being retained each time, the ammonia was frozen several times in liquid air, all uncondensed gases were pumped off, and during the periods of liquefaction some vapour was allowed to escape. Finally a product was obtained which contained less than 1 in 100,000 of incondensable gases and less than 0.003 per cent, of water. The boiling-point of this ammonia by the static method was –33.354° C.; by the ordinary method, using a Beckmann thermometer, it was -33.341° C.

Vapour pressures of liquid ammonia

Temperature, °C.	-78.0	-64.0	-51.0	-48.0	-44.0	-40.0	-38.87	-33.354
Pressure (mm.)	44.0	125.29	289.43	345.03	432.98	538.58	571.77	760.0
t° C.	-30.0	-25.0	-20.0	-15.0	-10.0	-5.0	0.0	5.0
p (mm.)	897.2	1137.4	1427.0	1772.8	2181.6	2661.5	3220.8	3867.9
t° C.	10.0	15.0	20.0	25.0	30.0	40.0	45.0	50.0
p (mm.)	4612.1	5462.6	6428.2	7520.5	8750.9	1166.4	13362.3	15245.4
t° C.	55.0	60.0	65.0	70.0
p (mm.)	17325.3	19606.0	22108.5	24836.8

These results and others can be expressed by the interpolation formulae

log₁₀(p) = 30.256818 – 1914.0569/T – 8.4598324 log(T) + 2.39309×10^-3T + 2.955214×10^-6T². (1)
log₁₀(p) = 12.463400 – 1468.6068/T – 0.01638646T + 2.403276×10^-5T² - 1.168708×10^-8T². (2)

These and the preceding data have much value in industry, since liquid ammonia is frequently used as a refrigerant. It is a convenient working substance for reversed heat engines, the latent heat of vaporisation and expansion through coils being absorbed by the brine of cold-storage rooms, etc. The gas, when used commercially, is dried with quicklime, and is kept in gasometers over petroleum oil; it is usually compressed in two stages, first to 3 or 4 atmospheres, then to liquefy at about 8 atmospheres.

Critical Constants

The critical temperature is variously given as 130°, 131°, 132.3°, and 132.9° C., the critical density as 0.2362, the critical volume as 4.234, and the critical pressure as 115, 113, and 112.3 atmospheres respectively.

Surface Tension σ.

Temperature, ° C.	11.10	34.05	58.98
σ	23.38	18.05	12.95

Adsorption by Charcoal

At ordinary temperatures 1 gram of charcoal adsorbs nearly 200 c.c. of ammonia. The amount adsorbed at atmospheric pressure increases with fall of temperature. At each temperature the volume "v" varies with the pressure according to the equation

v=ap^1/n,

in which a and n are constants. This equation expresses the behaviour of ammonia from -20° to +200° C., and is known as the adsorption isotherm. The heat of adsorption, 494 Cals. per gram of ammonia, is nearly twice as great as the latent heat of condensation to liquid.

Latent Heat

This is usually taken as 5.000 Cals. per mol. This is higher than the value obtained by the differentiation of equation (1) above, which gives 3.950 Cals.

The boiling-point of ammonia, considered as the first member of the series NH₃, PH₃, etc., is abnormally high, just as that of OH₂ is high in relation to those of SH₂, etc. This as well as other peculiarities in the properties of liquid ammonia, which it shares with water, are in both cases attributed to association of the liquids.

The value of the Despretz-Young-Trouton ratio, namely, the latent heat of vaporisation divided by normal boiling-point on the absolute scale, may be calculated from the equations of Cragoe, Meyers, and Taylor. It is 19.7. This is not far from the normal value, 20 to 22. It has been shown by Nernst that this constant is really a function of the temperature, and that in the case of many gases

Q/T_b = 9.5 log(T_b) - 0.007T_b.

Thus, at a boiling-point of 239.65, the normal constant will be 21. A slight degree of association is thus indicated.

The temperature coefficients of molar surface energy are 1.80 and 1.79 (vide supra). The corresponding value for a normal liquid being 2.12, the coefficient of association is 1.27.

The critical density, 0.2362, is 4.211 times the theoretical density, as deduced by a strict application of the gas laws to the critical conditions. The ratio in the case of normal liquids is about 3.6.

A comparison between liquid ammonia and water in respect of physical properties, and its action as-a solvent, reveals many interesting similarities. These are summarised in the following table:

Comparison of ammonia with water

	NH3	H2O
Melting-point, ° C.	-77	0
Boiling-point, ° C.	-38.5	+100
Density at 20° C.	0.61	0.9982
TC	273 +131 abs.	273+374 abs.
PC	113 (atm.)	218 (atm.)
Molar elevation of B.P. (100 grams)	3.4 (lowest known)	5.2
Heat of fusion	1.838 Cals.	1.440 Cals.
Dielectric constant	c. 21-23	81.7
Surface tension	25	28
Viscosity	2.6×10-3 (at its b.p.)	10.6×10-3 (at 18°)
Factor of association (from surface tension)	1.27	3.0
Refractive index	1.33	1.33

It will be observed that the fluidity is much greater (about four times) than that of water, and the surface tension appreciably lower, both being measured at ordinary temperatures. The various physical data lead to the conclusion that ammonia, like water, is probably associated in the liquid state. Like other pure liquids, ammonia is capable of conducting the electric current: its conductivity, although very small, has a definite limit, below which it is not diminished by the highest possible purification. As in the case of water, this is attributed to a very slight electrolytic dissociation:

NH₃ ⇔ H^• + NH'₂
and
NH₃+H^• ⇔ NH^•₄.

The dielectric constant is high, but far inferior to that of water. The refractive index is nearly the same as that of water. Ammonia also resembles water in showing some deep absorption bands in the infra-red between the wave-lengths 9μ, and 13μ. This is probably connected with its high ionising power as a solvent.

Liquid ammonia is one of the most comprehensive solvents known. It dissolves non-metals, iodine, sulphur, and phosphorus, also many halides and sulphides of the non-metals which are insoluble in water or are hydrolysed by it. The metals of the alkalies and alkaline earths are freely soluble, and the solutions are good conductors of electricity. The solubility of potassium in 100 grams of liquid ammonia is 48.75 at 0° C., that of sodium 23.2 at 0°, 24.5 at -30°, and 22.0 at 22° C. At such high concentrations the ammonia pressures are greatly reduced and may lie far below one atmosphere at ordinary temperatures. In the case of lithium two liquid phases are formed, having an ammonia pressure of 540 mm. The saturated solution, having an ammonia pressure of 3.61 mm., contains 11.4 grams of lithium in 100 grams of ammonia. Ammonia dissolves many typical aqueous electrolytes, such as salts. The following are insoluble: oxides, hydroxides, and sulphides of the metals; also fluorides, sulphates, sulphites, phosphates, arsenates, carbonates, and oxalates. The solubilities of other salts, especially of ammonium salts, is often high. That of ammonium nitrate, which is zero at -80° C., increases rapidly with rise of temperature.

t° C	-80	-30	0	+33.3
Grams of NH4NO3 in 100 grams of NH3	0	224	291	398

Dry ammonium thiocyanate cooled by ice absorbs 45 per cent, by weight of ammonia, which, however, evolves it again on heating. Such solutions have been used in small refrigerators, working somewhat on the principle of the Carre freezing machine.

Salts of the following acid radicals usually dissolve in liquid ammonia, namely, nitrates, iodides, bromides and chlorides, cyanides and thiocyanates. Other soluble substances are amides, such as KNH₂, non-electrolytes such as Hg(CN)₂, and various complexes such as K₂Hg(CN)₄, which often have the same composition as the corresponding complex salts which are formed in water. Many organic compounds will dissolve, among which may be mentioned amides, nitro-compounds, the free alkyl ammonium bases, their hydroxides, and salts.

Ammonia has been used as a solvent in the determination of molar weights. The molar elevation of the boiling-point (100 grams solvent), calculated from the thermal constants, is 3.4. This agrees with the experimental elevation in the case of some solutes, i.e. those which have a normal molar weight. As in the case of aqueous solutions, however, the molar weight found often differs from the formula weight. When the constant, or molar elevation of the boiling-point, calculated from the experimental elevation on the assumption of simple molecules, is higher than the theoretical constant, the solute is regarded as dissociated, as in the case of salts and some organic compounds, or else it is assumed to be associated with solvent molecules to give an ammine.

The molar weight of the solute may also be determined from the lowering of the vapour pressure. The molar weight of sodium found thus decreases nearly as a linear function of the concentration and does not reach a definite limit. It is only possible to state that the limit will be below 23.

The cryoscopic constant of ammonia, as experimentally determined, is about 14, whereas the value calculated from the thermal constants is 9.4.

The ions which are present in these electrolytically conducting solutions react and give precipitates, etc., which are sometimes of an unusual kind and would be difficult to produce in any other solvent. Dissolved metals may interact to give alloys; thus, in the case of sodium and lead, an alloy of composition represented by NaPb_2.25 is precipitated. The lead is the negative constituent, as it is precipitated on the anode during electrolysis. Sodium and tellurium give insoluble Na₂Te and soluble Na₂Te₂. Simple and complex amides and nitrides may be formed by double decomposition, e.g.:

3HgI₂+6KNH₂ = Hg₃N₂+6KI+4NH₃,
AgNO₃+KNH₂ = AgNH₂+KNO₃,
PbI₂+2KNH₂ = PbNH+2KI+NH₃.

Aluminium iodide with potassamide gives first Al(NH₂)₃.AlI₃, then Al(NH₂)₃.Al(NH₂)₂I.NH₃. Cadmium thiocyanate, with the same reagent, gives Cd(NH₂)₂, a white powder which acts violently on water and explodes when heated. Silver amide with potassamide gives crystals containing ammonia:

AgNH₂+KNH₂ = AgNHK.NH₃.

Magnesium dissolves in a solution of ammonium chloride in liquid ammonia. Here the ammonium ion plays the part of the hydrion in acid aqueous solutions, and is displaced by the magnesium:

Mg+2NH₄Cl = MgCl₂+H₂+2NH₃.

The strongly basic alkyl ammonium hydroxides displace the ammonia from its salts, as they do in aqueous solution:

(CH₃)₄NOH+NH₄Cl = (CH₃)₄NCl+NH₃+H₂O.

Since in ammonia the - NH₂ group may be regarded as taking the place of the - OH in water, acid amides in ammonia correspond to carboxylic acids in water, and the former have therefore been named "ammono-acids" by Franklin. They are dibasic acids, capable of ionising in two stages into RCONH'+H^• and RCON''+2H^••. The ionisation is in some cases considerable, as shown by the conductivities. Typical reactions of the ammono-acids are:

Mg+CH₃CONH₂ = CH₃CONMg+H₂,
CH₃CONH₂+(CH₃)4NOH = CH₃CONH.N(CH₃)₄+H₂O,
AgNH₂+CH₃CONH₂+NH₃ = CH₃CONHAg.2NH₃.

In water, acid amides do not dissociate but add on the hydrion of acids, and therefore behave as (weak) bases. The amides of bases have been called by Franklin " ammono-bases." These will react with ammono-acids thus:

CH₃CONH₂+KNH₂ = CH₃CONHK+NH₃,
or
CH₃CONH'+H^•+K^•+NH'₂ = CH₃CONH'+K^•+NH₃.

As in the case of the neutralisation of an acid by an alkali in water, the ionic part of the reaction consists in the formation of undissociated solvent:

H^•+NH'₂ = NH₃.

Just as salts of weak bases are hydrolysed by the substitution of OH for halogen, etc., so they may be " ammonolysed " by the substitution of NH₂:

HgCl₂+NH'₂+NH^•₄ = HgNH₂Cl +NH₄Cl.

The amido-chloride is insoluble in ammonia as it is in water.

Ammonolysis is, however, much less common than hydrolysis. Most salts dissolved give clear solutions in ammonia, while ammonolysis would give precipitates.

The ammono-salts of the alkali metals are also obtained by reactions in liquid ammonia. Thus, dipotassium ammono-sodiate, Na(NH₂)₃K₂, is prepared by the action of sodium on potassamide, both dissolved in ammonia, in the presence of platinum black.

Amides and nitrides of the non-metals are often formed by the ammonolysis of the chloranhydrides and sulphides in liquid ammonia:

3TeCl₄+16NH₃ = Te₃N₄+12NH₄Cl,
B₂S₃+6NH₃ = B₂(NH₃)₃+3NH₄SH,
PCl₃+5NH₃ = P(NH)(NH₂)+3NH₄Cl,
SiCl₄+8NH₃ = Si(NH₂)₄+4NH₄Cl,
SiS₂+4NH₃ = Si(NH₂)₂+2NH₄SH.

The electrical conductivity of pure ammonia is very low, and in this case the difficulty of avoiding traces of water are very great. Hence the older values of this constant are higher than the more recent, namely, 0.1×10^-7 reciprocal ohms.

Besides the ions H^•, NH^•₄, and NH'₂, the anions NH'' and N''' may possibly be present. Three inflections have been found on the curve of anodic polarisation against current, which may correspond to the discharge of these ions. In accordance with its character as a dissociating solvent, liquid ammonia has the relatively high dielectric constant of 21 to 23, which is inferior only to those of some alcohols and nitriles, e.g. benzonitrile, 26.0, and to that of water, 81.45. In the case of ammonia, as in that of water, there are two types of solutions: those of salts which have a limiting conductivity at limiting dilution, but do not give a dissociation constant, i.e. do not obey the dilution law; and those of organic compounds of a polar type, which do not give a limiting equivalent conductivity but do obey the dilution law. As an example of the first class we may take KNO₃. The equivalent conductivities, λ, and degrees of dissociation, a, at -33° C. and various dilutions, V (litres per mol.), are as follows:

V=	324	1001	2514	6162	23060	69820	∞
λ=	192.7	245.0	282.7	309.9	330.1	338.6	339 = λ0
α	0.567	0.720	0.831	0.912	0.972	0.995	1.00

The equivalent conductivities are higher than in water at corresponding and at infinite dilutions, but the degrees of dissociation at corresponding dilutions are lower. The following equivalent conductivities are additively made up of a term due to each ion:

Salt	KNO3	KBr	NaNO3	NaBr	NH4NO3	NH4Br	LiNO3	AgNO3
λ0=	339	340	301	302	302	303	283	287

Since the transport numbers of salts in ammonia have been determined, the ionic conductivities can be calculated and prove to be about three times the conductivities in water.

Kation	Li	Na	K	NH4	Ag
L0 in H2O (18° C.)	33.3	43.4	64.5	64.7	54.0
L0 in NH3 (at b.p.)	112	130	168	131	116

Anion	Cl	Br	I	NO3
L0 in H2O (18° C.)	65.5	67.7	66.6	61.8
L0 in NH3 (at b.p.)	179	172	171	171

The differences between individual ions, especially anions, are much less marked in ammonia than in water. The higher conductivities are no doubt connected with the lower viscosity (or higher fluidity) of ammonia, since the resistance to the motion of the ions is some function of the viscosity. That of water at 18° C. is 10.63×10^-3, while that of ammonia at its boiling-point is 2.56×10^-3.

In the case of the slowest ions, the ratios of ionic conductivity to fluidity are about 0.3 in both solvents.

In very dilute solutions of salts, and in rather dilute solutions of other compounds, a mass action constant may be calculated, but in order to express the behaviour in more concentrated solutions other constants must be introduced, as in the equation of Kraus and Bray:



Nitromethane, thiobenzamide, orthonitrophenol, methylnitramine, trinitrobenzene, trinitraniline are among the organic compounds which are electrolytes in ammonia and obey the law of mass action moderately well. The following are typical constants:

	104k	m.	D.	λ0
Potassamide	1.20	1.18	0.095	301
Trinitraniline	30.0	0.73	0.38	234

Liquid ammonia is a favourable medium for the preparation of free "ammonium" and the alkyl substituted ammonium radicals. The radical N(CH₃)₄ has been precipitated in the free state at the kathode during the electrolysis of the salt dissolved in liquid ammonia, in which it is soluble, giving a solution somewhat similar to that of potassium. When compressed it is a good conductor of electricity. The radical HgCH₃ can be freed from HgCH₃Cl by washing with liquid ammonia. It decomposes at ordinary temperatures into Hg(CH₃)₂ and Hg.

Solutions of the alkali metals are blue when dilute; they assume a metallic lustre and become reflecting when more concentrated. There is little chemical action; the amides are only formed to a slight extent, so that it must be supposed that the metal is dissolved as such, and that the high conductivity is not due to ordinary ions. During electrolysis there are none of the usual appearances at the electrodes, such as the separation of solids or gases. The conduction in this respect resembles that of a metal. The dissolved metal is associated with the kation, since it may be completely transferred into the katholyte by electrolysis. The negative carrier is probably the electron. The equivalent conductivity (referred in the usual manner to the solute, e.g. dissolved sodium) is high; it reaches a minimum at a dilution of between 20 and 30, and then increases greatly with increase of concentration:

V	0.5047	1.038	13.86	30.4	690.1	37880
λ	82490	3228	478.3	478.5	869.4	1034

The specific conductivity =λ/1000V obviously increases still more rapidly with increasing concentration, and the atomic conductance approaches that of the metal itself, as is seen from the results in still more concentrated solutions:

V	1.674	0.9265	0.5099	0.3230	0.1081
x	1.298	5.988	148.3	714.0	5047.0

The atomic conductance of a saturated solution is 1.1×10⁶ mhos, that of the metal 5.05×10⁶ mhos.

The transport number of the anion has been investigated by the method of concentration cells. The e.m.f. of these at concentrations from 0.010 to 0.870 shows an enormous increase in the velocity of the negative carrier, which probably consists to a greater and greater extent of electrons, the conductivity at high concentrations being due to the dissociation:

M ⇔ M⁺+e^-

In more dilute solutions some amide may be formed, which then dissociates into ions which have velocities of the usual order:

M+NH₃ = MNH₂+½H₂,
MNH₂ ⇔ M⁺+NH₂^-1.