equivalent to 1.137 grams I]. Fill both globes with dry hydrogen, in the manner described under I. for hydriodic acid: seal, and heat to 440° for 4 hours. Then cool quickly, and determine the ratio of free hydrogen to total hydrogen as before.

Now calculate the ratio of hydriodic acid formed to hydriodic acid possible, in each case; from the equation

Your experiments ought to shew that as a decreases, à increases; or, in other words, that, other conditions being the same the dissociation-limit is lessened when either of the products of dissociation is present in excess.

Exp. 2. Dissociation of ammonium carbamate. Read Horstmann, Annalen der Chem. u. Pharm. 187, 48.

I. Dissociation of ammonium carbamate in vacuo. Procure a tube, closed at one end, about a metre in length and about 15 mm. in diameter. The tube is provided with a millimetre scale in terms of which it is accurately calibrated. Having thoroughly cleaned and dried this tube, lead into it perfectly dry ammonia and carbon dioxide gases. Make the delivery tubes so long that the gases first come into contact with each other near to the closed end of the tube. White crystalline plates of ammonium carbamate CO(NH)(ONH1) form on the sides of the experimental tube. When this formation has gone on for some time, fill the tube very carefully with perfectly dry (but not warm) mercury, and invert it in a trough of mercury. It will be found that the height of the mercury in the experimental tube is now less than the height of the barometer. This is due to the fact that at ordinary temperatures, i.e. 17o-20o, the carbamate undergoes a notable amount of dissociation into carbon dioxide and ammonia.

The difference between the barometric height and the height of mercury in the experimental tube gives the dissociation-pressure for the temperature considered.

The following are some of the values obtained by Horst

mann for the dissociation-pressures of the carbamate at ordinary summer temperatures ;—

It is evident from these numbers that the dissociation increases with the temperature.

The student should perform three or four experiments in the manner described to verify this result.

II. Dissociation of ammonium carbamate in presence of excess of either of its products of dissociation.

It can be theoretically proved (see Horstmann loc. cit.) that although the dissociation-pressure of ammonium carbamate in presence of excess of either of its products of dissociation ought to be less than what it would be in vacuo at the same temperature, yet an excess of ammonia ought to diminish the dissociation-pressure to a greater extent than an equal excess of carbon dioxide.

Prove this experimentally. The apparatus required is the experimental tube already described and a small gasometer similar to those used in nitrogen-determinations; the gasometer is calibrated with the same mass of mercury as was used for the calibration of the experimental tube. The perfectly dried gas (either ammonia or carbon dioxide as the case may be) is brought into the gasometer and its volume is read off. The end of the delivery tube of the gasometer is then brought under the mouth of the experimental tube, and gas is forced into this by pouring mercury into the funnel tube. When sufficient gas has passed into the experimental tube the volume of residual gas in the gasometer is read off. We have now all the data necessary for calculating the partial pressure, P, which the gas introduced into the experimental tube would exert there.

The difference between the total pressure in the tube when this has become constant and the partial pressure P gives us the dissociation-pressure (p,) of the carbamate under the conditions of the experiment.

It is scarcely necessary to add, that the experiments should be conducted in a room the temperature of which

remains or can be maintained as nearly as possible constant for intervals of 3 or 4 hours.

Appended are some of the results obtained by Horst

In these tables p denotes the dissociation-pressure of the

carbamate in vacuo.

Four or five experiments should be made with different partial pressures in the case of each gas.

If the results in each case are plotted out in such a way

Ρ

obtained which are independent of the various temperatures at which the experiments have been conducted. In accordance with what has been said above, it will be found that the curve for ammonia is steeper than that for carbon dioxide.

CHAPTER III.

RELATIVE AFFINITIES OF ACIDS.

WHEN equivalent masses of two acids and a base are allowed to interact in dilute aqueous solution, the base divides itself between the acids in a definite ratio; this ratio expresses the relative affinities of the acids for the base. The values of the relative affinities of acids quantitatively condition many, if not all, chemical changes which are brought about by the acids. (s. Pattison Muir's Principles of Chemistry, Book II. Chap. III., where references are also given to original memoirs of importance.)

Exp. 1. Thermal methods of measuring relative affinities of acids. Read Thomsen, Thermochemische Untersuchungen, 1, 97-126.

When equivalent masses of sulphuric acid and caustic soda interact in dilute aqueous solution, a gram-units of heat are produced; when equivalent masses of nitric acid and soda interact, b gram units of heat are produced. When equivalent masses of the two acids and soda interact, either the whole of the soda forms sulphate, in which case a units of heat are produced; or the whole of it forms nitrate, in which case b units of heat are produced; or both sulphate and nitrate are formed, in which case the quantity of heat produced is different from either a or b.

Measurement of the heat actually produced will furnish data from which conclusions may be drawn regarding the distribution of the soda between the two acids.

When equivalent masses of sodium sulphate and nitric acid interact in dilute aqueous solution, either no change

occurs, in which case no thermal disturbance is noticed; or the whole of the sulphate is decomposed, in which case a thermal change, equal in amount to the difference between a and b, occurs; or a portion of the sulphate is decomposed, in which case the amount of thermal change is less than the whole difference between a and b.

Supposing that b<a, and that when sodium sulphate and nitric acid interact x grams-units of heat disappear, a being less than b-a; then, if no other changes occur than those taken into account, we may say that of the total mass

b

a

of sodium sulphate has been decomposed; hence we might calculate the distribution of the soda between the two acids, and so find the relative affinities of the acids. Thus, Thomsen found a = 31,378, b = 27,234 (b − a = — 4144), x = — - 3504; hence, assuming that no changes occur except those represented by

a and b,

- 3504
4144

845 of the total sodium sulphate has been

decomposed by the nitric acid.

But if, when sodium sulphate and nitric acid interact, the whole of the sulphate is not decomposed, the solution must contain sodium sulphate and nitrate, and also sulphuric and nitric acids; either or both of the acids may interact with either or both of the salts, or the acids may interact with each other, and these changes must be accompanied by production or disappearance of heat. Thomsen has found that the only one of the possible reactions accompanied by more than a very small thermal change is that between the sulphuric acid and the sodium sulphate. Hence it is possible to find by thermal methods the approximate distribution of soda between two acids when the three bodies interact in dilute solution in equivalent masses.

The data to be determined are the thermal values of the following chemical changes;— (1) [H3SO1Aq, 2NaOHAq]; (2) [2HNO3Aq, 2NaOHAq]; (3) [Na3SO1Aq, 2HNO3Aq]; (4) [Na'SO Aq, nH2SO1Aq].

Thomsen's solution of the problem is given to shew how the calculations are made.

(1)=31,378; (2) = 27,234, .. (2) -(1)=-4144;