CHAPTER VII

«PHYSICAL OR HETEROGENEOUS EQUILIBRIUM.—THE COLLOIDAL CONDITION»

The law governing ‹physical› or ‹heterogeneous equilibrium› applies to all cases where, at a constant temperature, one and the same chemical substance is present in two or more physical conditions, or "phases," in contact and in equilibrium with each other. We have, for instance, the common case of a liquid, say water, in contact with its vapor, or the liquid in contact with its solid phase (ice) and its vapor; or, we may have a gas, say oxygen, in contact with its solution in some solvent like water. We may have a solid, like cane sugar, in contact with its solution. We may also have a substance like bromine, which is soluble both in chloroform and in water, present in both solutions at the same time, the two solutions being in contact but immiscible. These cases represent the most common types of systems to which the law of physical equilibrium may be applied, although the list has by no means been exhausted. ‹The law of physical or heterogeneous equilibrium states that when one and the same chemical compound is present in two physical states or phases›, as expressed in the equation S_{1} ⇄ S_{2}, then ‹when equilibrium is reached›, at a given temperature, ‹the ratio of the concentrations of the substance in the two phases is some constant number›:

[S_{1}] : [S_{2}] = «k.»

The bracketed symbols denote concentrations.

It should be noted that the condition of equilibrium is independent of the total quantity[227] of substance present in either phase or in both phases; that is, provided the ‹ratio› of the ‹concentrations› is maintained constant at a given temperature, the ‹quantity of substance present in both phases or in either phase is variable›. For instance, the condition of equilibrium between water and [p119] water vapor is independent of the quantity of water or of water vapor present: in a closed liter bottle containing water and water vapor, the ratio of the concentrations is maintained, irrespective of the question whether the bottle contains 10 c.c. of water and 990 c.c. of vapor or 990 c.c. of water and 10 c.c. of vapor.

The law is one of experience; instances of its application are given below. Its probable theoretical significance may be explained mechanically with the aid of the kinetic theory of gases and solutions, as follows: If chloroform is added to a solution of bromine in water, the chloroform takes up part of the bromine and, if the mixture is vigorously shaken, a condition of equilibrium and a definite distribution of bromine between the two solvents will result (‹exp.›[228]). Now, if one imagines a liter of the aqueous solution to contain ‹one mole› of bromine at some given temperature and to cover a liter of chloroform, the whole system being left to itself, then all the conditions affecting the migration of the bromine into the chloroform will be definite ones—the concentration of the bromine, the temperature, the surface between the two solvents—and bromine will pass from the aqueous solution into the chloroform solution at a definite speed. We may call the ‹quantity› (in moles) of bromine which would enter the chloroform in one minute, if the concentration of the bromine in the water were kept constant (one mole) throughout the minute, the ‹velocity of migration› of the bromine—this velocity, like chemical velocity, representing a quantity, not a distance. The velocity being a definite one under these conditions, we have ‹v›_{1} = ‹k›_{1}. Now, if all the conditions are left unaltered, except that the concentration of the bromine is changed, say kept at one-hundredth its original value, then only one one-hundredth as many molecules of bromine as in the first case will come into contact with the chloroform surface in unit time. The chances for migration are one one-hundredth as great, and the quantity entering the chloroform in unit time—the velocity of the [p120] change—will be one one-hundredth of the original velocity. In general, ‹the velocity will be proportional to the concentration of the bromine› [Br]_{aq.} ‹in the water at any moment› and to the characteristic constant ‹k›_{1}.

‹v›_{1} = [Br]_{aq.} × ‹k›_{1}.

On the other hand, if a solution of bromine in chloroform is covered with water, bromine enters the water (‹exp.›).[229] We would find, by the method of analysis used before, and for the same conditions, that the velocity of migration, ‹v›_{2}, of the bromine into the water is also proportional to a characteristic constant, ‹k›_{2}, and to the concentration of the bromine in the chloroform [Br]_{ch.}. We have, therefore ‹v›_{2} = [Br]_{ch.} × ‹k›_{2}.

Equilibrium between the two solutions will be reached when

‹v›_{1} = ‹v›_{2} or [Br]_{aq.} × ‹k›_{1} = [Br]_{ch.} × ‹k›_{2},

from which follows that[230] for the condition of equilibrium

[Br]_{aq.} / [Br]_{ch.} = ‹k›_{2} / ‹k›_{1} = ‹k›.

«Applications of the Law of Physical Equilibrium.»—(1) The law of physical equilibrium may be applied first to the case of a liquid, say chloroform, in contact with its vapor. For the condition of equilibrium at a fixed temperature [CHCl_{3}]_{vap.} : [CHCl_{3}]_{liq.} = ‹k›.

Now, at a fixed temperature, a pure liquid has a fixed concentration, its specific gravity being a definite one. Hence, for a fixed temperature, the second term of the constant ratio being definite, the first term, [CHCl_{3}]_{vap.}, representing the concentration of the vapor, must also have a fixed, constant value for the condition of equilibrium, ‹i.e.› when the space above the liquid is saturated with its vapor. This is in agreement with well-known facts. The concentration of the vapor is usually expressed in terms of its [p121] pressure, and is called the ‹vapor pressure› or ‹vapor tension› of the liquid at the temperature in question. Tables giving the definite vapor pressures of important liquids at the various fixed temperatures are in common use.

(2) For oxygen in equilibrium with its saturated solution, say in water, at a fixed temperature, we have, according to the law, [O_{2}]_{gas} : [O_{2}]_{solut.} = ‹k›.

If the oxygen is under a given pressure at a definite temperature, its concentration [O_{2}]_{gas} is fixed, and consequently the second term, [O_{2}]_{solut.}, of the ratio, the concentration of the dissolved oxygen, or ‹its solubility, must also be definite›, ‹i.e.› oxygen must have a definite solubility in water at a given temperature under a given pressure. If the pressure on the gas is doubled, its concentration is doubled and, to maintain the constant ratio, its solubility must also be doubled—which is in agreement with the facts (Henry's law). If air of the same pressure is taken, in place of pure oxygen, then the concentration of the oxygen (first term of the above ratio) is only about one-fifth as great as for the pure gas, and the water must be saturated with oxygen when it has taken up only one-fifth (second term of the ratio) as much as it would from the pure gas (Dalton's law): ‹Each gas in a mixture is soluble in proportion to its own partial pressure› or ‹concentration›.

(3) For sugar in equilibrium with its solution in water, ‹i.e.› in contact with its saturated solution, at a given temperature, we should have

[C_{12}H_{22}O_{11}]_{aq.} : [C_{12}H_{22}O_{11}]_{solid} = ‹k›.

Since a pure solid like sugar at a given temperature has a definite specific gravity, the concentration of the sugar in the solid condition should also be a definite one. Consequently, according to the law under discussion, the first term, [C_{12}H_{22}O_{11}]_{aq.}, of the ratio, the concentration of the sugar in its saturated solution in contact with the solid phase, must also have a definite value. This is in agreement with fact, the concentration of the sugar in the saturated solution, termed its solubility, being a definite one for a given temperature. (See below, p. 123, in regard to the solubility of fine powders.)

«Supersaturated Solutions.»— It is well known, however, that by dissolving a substance, such as sugar, in hot water and carefully [p122] cooling the solution, we may obtain a solution of sugar containing much more sugar in unit volume than is represented by its solubility at the lower temperature. The concentration of the sugar at the temperature in question, instead of having the definite value represented by [C_{12}H_{22}O_{11}]_{aq.}, can easily have a value several times as large. This phenomenon ‹is not at variance with the law of physical equilibrium›, inasmuch as the law states that, when a given compound is present in ‹two› physical states or phases, then a constant ratio between the concentrations of the substance in the two phases is established when equilibrium is reached at a given temperature. In the solution prepared as described, we have the substance present only in one phase, and we have what may be called a metastable condition, as long as the second phase is not introduced. As soon as a minute crystal of the solid phase is added or is formed in the solution, change immediately ensues, and the excess of solid is deposited. If the mixture be kept perfectly quiet, the excess will in most cases be deposited on the ‹surfaces› of the added crystal, which thereby grows larger (rock-candy manufacture).

EXP. Supersaturated solutions of sodium sulphate and sodium thiosulphate, into which crystals of the salts are dropped, show how the crystal starts crystallization. The crystals develop as branches from the crystal first introduced and from the new crystals formed.

This phenomenon of ‹supersaturation› is one which analytical chemists must always take into consideration. Tests which involve the precipitation of substances that are merely difficultly soluble, rather than exceedingly insoluble, or of substances present only in very small quantities, may well lead to entirely wrong conclusions, if precautions are not taken against the possibility of the failure of a precipitate to appear as a consequence of persistent supersaturation. For instance, a common test for the presence of potassium salts consists in the precipitation of ‹potassium acid tartrate› by the addition of tartaric acid to the solution of a potassium salt (‹exp.›). The tartrate is somewhat soluble and tends to form supersaturated solutions; if we proceed without due regard for this phenomenon, we may readily have a quantity of potassium salt present and fail to obtain the test for it. Simply mixing tartaric acid and potassium chloride solutions (‹exp.›) may fail to [p123] give any precipitate, and if the test is thrown away and potassium reported absent, a glaring blunder is committed. To insure against the error of supersaturation, we try to start crystallization by the common devices of shaking the solution or "scratching" the walls of the vessel, the object being to facilitate the formation of the first crystal. The surest method is to ‹inoculate› a small portion of the mixture with a ‹minute› crystal of the substance we expect to be formed. If no precipitate results in a short time, the solution is not supersaturated—it may be too dilute and may require further concentration, but the error of supersaturation has been excluded.

The relation between supersaturated solutions and crystals brings out sharply the fact that physical equilibrium is essentially a condition of equilibrium between the substance at the ‹surface› of the solid and the substance in its dissolved state. In terms of the molecular theory, equilibrium is established when the molecules of the crystal surface dissolve as rapidly as molecules from the solution are deposited on the surface. If the concentration of the dissolved molecules is reduced below the point required for equilibrium, the velocity of deposition is diminished. The velocity of solution will then be greater than the velocity of deposition and ‹solution› will result. The reversed relations hold when the concentration of the solution is greater than that demanded by equilibrium: the velocity of deposition will be the greater than that of solution and ‹precipitation› follows.[231]

«Solubility of Fine Powders.»—Consideration of the surface forces, acting between crystals and the liquids wetting them, led to the interesting prediction[232] that the more minute crystals of a given specimen would not only dissolve more rapidly, on account of the larger surface exposed, but would also be ‹more soluble› than the larger crystals, and for the same reason. Surface tension always tends to produce the smallest possible free surface of a liquid, and the free surface between a liquid and a given weight of solid material in a fine powder is much larger than between the liquid and the same weight in larger crystals. The surface tension [p124] will therefore tend to convert the smaller crystals into larger ones, and it can do so only by means of a greater degree of solubility of the former. This prediction has now been fully verified by experiments on the solubility of barium sulphate and of gypsum.[233] The solubility of barium sulphate in a very fine powder (with an average diameter of 10^{−4} mm.) is almost twice as great as the solubility of a coarser material (18E−4 mm. average diameter).

The application of these relations to analysis is as follows: If a crystalline precipitate is in contact with a solvent, ‹e.g.› if barium sulphate is in contact with the liquid from which it has been precipitated, then this liquid must be continually in a state of change, not of equilibrium, with respect to the solution and the deposited barium sulphate. The more minute crystals, being a little more soluble than the larger ones, will supersaturate the solution in respect to the larger crystals and the excess will be deposited on these larger crystals and make them grow still larger. This deposition will make the solution unsaturated with respect to the smaller crystals and more of these will dissolve. The process is obviously a continuous one, and must lead in time to the disappearance of the minute crystals and the growth of the larger ones. That is a result which analysts aim to attain,—which in quantitative work it is in fact necessary to attain, since the more minute crystals are likely to pass through filters and be lost in the analysis. The views expressed, and the experimental confirmation of the conclusion reached, form the theory of what is called the "digesting" of precipitates before they are brought on filters. It is clear that every condition facilitating contact between solvent and solid will accelerate the desired change and continuous ‹stirring› is therefore desirable. Heating is, as a rule, also to be desired for very insoluble precipitates, as it will, in the majority of cases, facilitate the solution of the undesirable, finer crystals.

In conclusion, these considerations will also indicate the precautions to be observed in the precipitation of difficultly soluble substances. If this is not properly carried out, endless trouble in [p125] the analytical laboratory results. Except when heating is, for some special reason, undesirable—as inducing a chemical change (like hydrolysis) that is not wanted—solutions are brought to the boiling-point, and ‹the precipitant added drop by drop›, in order not to supersaturate the solution too strongly. The solution is thus allowed time to deposit its excess as far as possible on the ‹first crystals formed›, which it will do rather than to form new, minute ones (see supersaturation, p. 121). ‹Constant stirring› is prescribed in order to bring older crystals, as far as possible, into contact with all parts of the slightly supersaturated solution. After the precipitation is complete, it is usually desirable to allow the mixture to "digest" for some time, for the reasons given above—stirring and a high temperature during the process being desirable. (See further Chap. VIII, p. 147, ‹in regard to the use of an excess of precipitant›.)

THE COLLOIDAL CONDITION

When a difficultly soluble substance is formed in a solution beyond the point of saturation of the solution, the substance in question ‹separates› from the solution ‹in a new phase›, according to the principles just laid down. Ordinarily, if the substance is a solid, a ‹precipitate› is formed; if a gas, a gas escapes; if a liquid, a liquid separates out, which is immiscible with the solution in which it is formed. Occasionally, the condition of supersaturation which precedes the separation is somewhat persistent, but this resistance to the separation of the phase may be overcome by vigorous agitation of the solution or, as in the case of supersaturation with a crystallizable salt, by inoculation of the solution with a particle of the new phase (p. 123).

Under certain conditions, however, a difficultly soluble substance may be produced in a solution, in a concentration far beyond its solubility, ‹without the separation of a precipitate› (evolution of a gas or formation of a separate liquid) ‹and also without the formation of a supersaturated solution›. Thus, when hydrogen sulphide is passed into an aqueous solution of arsenious oxide, the liquid acquires the yellow to orange color of ‹arsenious sulphide› and becomes opalescent. But no ‹precipitate› is seen (‹exp.›), in spite of the fact that the sulphide is extremely insoluble and is formed practically quantitatively according to the equation As_{2}O_{3} + 3 H_{2}S ⇄ As_{2}S_{3} + 3 H_{2}O. The orange liquid passes through a [p126] filter unchanged (‹exp.›). But if some hydrochloric acid or a salt (‹e.g.› sodium chloride) solution is added to a portion of it, a heavy precipitate of arsenious sulphide is immediately produced (‹exp.›); its quantity is a good indication of the great amount of sulphide that is ‹not› precipitated before the addition of the acid or salt. If some pure arsenic sulphide (solid) is added to another portion of the orange liquid, in order to overcome any possible condition of supersaturation, the liquid is found to remain clear (but opalescent), excepting for the few particles of added sulphide (‹exp.›); even when it is vigorously shaken (‹exp.›), or allowed to stand for days, no precipitate is formed. We are therefore not dealing with a case of supersaturation.

A liquid, in which a very insoluble substance appears thus to be in solution far beyond its usual degree of solubility, and yet does not show at all the behavior of a supersaturated solution, is said to contain the substance in the «colloidal condition». After a few more instances of the colloidal condition have been presented, the significance of the condition and the meaning of the term used to designate it will be explained.

«Colloidal Gold.»—When a solution of gold chloride[234] is treated with a solution of stannous chloride which contains a little stannic chloride,[235] a purple-red, flocculent precipitate—the "purple of Cassius"—is formed; in extremely dilute solutions only a purple-red liquid[236] may be produced. If the precipitate is collected on a filter, washed with water and then treated, on the filter, with a few drops of ammonium hydroxide solution and some water (‹exp.›), it is seen to ‹dissolve› and a beautifully colored liquid, of purple-red or claret-red tint, is found to pass through the filter. In spite of the extreme insolubility of metallic gold, the red ammoniacal solution (as well as the first red precipitate) contains «metallic gold», in the colloidal condition, formed according to the equation[237] 2 Au^{3+} + 3 Sn^{2+} ⥂ 2 Au + 3 Sn^{4+}. By the careful [p127] reduction of gold chloride with phosphorus,[238] or with formaldehyde in dilute, slightly alkaline solution,[239] brilliant red liquids, containing metallic gold in the colloidal condition, may be prepared, which remain clear for months. The passage of an electric current, in the form of an arc, between two gold points under pure water produces similar red liquids[240] containing metallic gold.

«Colloidal Silver.»—Further, although silver is likewise an extremely insoluble metal, solid preparations of silver are known which, on treatment with water, form apparently perfectly clear (opalescent) liquids or solutions (‹exp.›).[241] From these silver is not deposited, even in the course of months. If a little of the black solid is heated on the lid of a porcelain crucible, the metal may be readily recognized by its white color and luster.[242]

«Colloidal Ferric Hydroxide.»—Further, if to a solution of ferric chloride an excess of ammonium hydroxide is added, the well-known, rust-red precipitate of very difficultly soluble ferric hydroxide is formed: FeCl_{3} + 3 NH_{4}OH ⥂ Fe(OH)_{3} ↓ + 3 NH_{4}Cl. The precipitate is so insoluble that it is a favorable form for precipitating the ferric-ion in quantitative analysis. If, however, a solution of ferric chloride be carefully neutralized with ammonium carbonate[243] and be then placed in a vessel (called a dialyzer), in such a way that it is separated, by an animal membrane or by parchment, from pure water, ferric hydroxide is obtained in an apparently soluble form. The ammonium chloride, as well as hydrochloric acid which is formed by the decomposition of the chloride by water (FeCl_{3} + 3 HOH ⇄ Fe(OH)_{3} + 3 HCl, see Chapter X), are found to pass through the membrane readily, while the ferric hydroxide produced does not pass through such [p128] membranes[244] and is retained in the dialyzer ‹without forming a precipitate›. The acid and salt pass through such membranes in either direction, indeed, but flow mainly from their solutions of higher concentration to those of lower concentration. The water on the outside of the dialyzer is, therefore, continuously renewed, in order to insure a concentration of these substances on the outside of the dialyzer lower than that within it. As the hydrochloric acid is thus removed through the membrane, the condition of equilibrium between the ferric chloride, water, ferric hydroxide and acid is continuously disturbed and, in the reversible reaction, expressed in the above equation, the action is carried more and more towards the right, and more and more ferric hydroxide is formed. The salt is, at last, practically all decomposed and a ‹clear red opalescent liquid, which contains ferric hydroxide in enormous excess beyond its solubility in water›, is left in the dialyzer.

Exactly as in the case of the other liquids discussed above, in which very insoluble substances appear to be in solution far beyond their usual degree of solubility, so the present liquid does not show the behavior of a supersaturated solution and it is ‹said to contain the ferric hydroxide in the› «colloidal condition».[245]

«Solution Theory of the Colloidal Condition.»—For many years the belief was prevalent among chemists that these liquids represent true solutions of difficultly soluble substances, in the form of soluble (colloidal) ‹modifications› of the substances. Like ordinary [p129] solutions, the liquids are found, indeed, to show a certain ‹osmotic pressure›, but, unlike ordinary solutions, the osmotic pressure is exceedingly small in proportion to the quantity of the substance present. Since the osmotic pressure is proportional to the number of molecules in unit volume (Chap. II), this observation proved the presence of a relatively very small number of molecules in solution, and chemists were led to assign, therefore, a relatively very great weight to each. Hence, the data led to the assumption that the colloidal modifications consist of molecules of very large, sometimes enormous molecular weight. From the fact that colloids are unable to traverse membranes, through which crystalloids readily pass, Graham had reached the same conclusion.

The following measurements of osmotic pressures may be given:[246]

Osmotic Substance. Concentration Pressure. Per cent. Cm. Hg. Gum arabic 1 6.9 Dextrin 1 16.6 As_{2}S_{3} 4 1.7 Fe(OH)_{3} 1.1 0.8 2.0 2.8 3.0 5.6 5.3 12.5 8.9 22.6

The last results, obtained with special care to exclude soluble impurities, are more reliable than the older results on gum arabic and dextrin. For the purpose of comparison it may be said that a 1% solution of cane sugar (molecular weight = 342) has an osmotic pressure of 53 cm. Hg at 18°, a 5% solution a pressure of 265 cm.

«The Suspension Theory of the Colloidal Condition.»—On the other hand, a number of chemists considered the colloidal liquids to represent, essentially, ‹suspensions of minute solid particles›[247] in an extreme state of subdivision, or, in some instances, perfect [p130] emulsions of difficultly miscible liquids.[248] Observations with the ultramicroscope[249] finally proved,[250] beyond question, the correctness of this theory of the colloidal condition. Thus, the colloidal "solutions" of gold are seen to contain minute solid particles of gold, the diameter of which varies[251] between (approximately) 60 µµ and 6 µµ and the color of which varies with their size. Still finer subdivisions, the diameter of whose particles cannot be measured, are also found to exist. Colloidal silver, platinum, arsenious sulphide and other colloidal metals and sulphides have also been shown, in the same way, to be suspensions of solid particles.

«The General Character and the Definition of the Colloidal Condition.»—Recent investigations have shown that the colloidal condition is possible, not only for a limited class of substances, but is, in general, possible for all substances.[252] Thus, even such an eminently crystallizable, readily soluble (in water) substance as sodium chloride may be obtained, under certain conditions, in colloidal suspension in benzene,[253] in which it is insoluble.

For our purposes it will be sufficient to define the colloidal condition, on the basis of these results, as the condition of an insoluble substance in which, as far as ordinary observation and the common methods for separation of heterogeneous phases (filtration, sedimentation, etc.) are concerned, ‹the substance appears to be present in a homogeneous clear solution›, whereas ‹in reality› it is present in a ‹heterogeneous mixture›. An extremely finely divided solid suspended in a liquid, or an emulsified liquid suspended in another liquid, are the most common types[254] of such mixtures. [p131]

«Relations to Analysis.»—Since the colloidal condition of insoluble substances interferes with the precipitation of the latter, and with their separation by filtration from the liquids in which they are suspended, and since the majority of the separations and tests of analytical chemistry depend on the successful formation of precipitates and on their successful separation by filtration, analytical chemistry is primarily[255] concerned with the colloidal condition as a condition that is to be ‹avoided as completely as possible› in order to escape error. In other fields of chemistry, notably in physiological chemistry and in some branches of technical chemistry, it is a source of effects so momentous and specific that a new branch of chemistry, the chemistry of colloids,[256] has grown out of the investigations of its relations and laws. The present discussion will be limited to the presentation of such of the fundamental facts concerning the colloidal condition as are of chief importance in analytical work.[257]

«Electrical Conditions of Colloids.»—One of the most important discoveries made on colloids is the observation that the suspended particles of a large class of colloids carry electrical charges, so that there is a potential difference between the particles and the liquid in which the suspension exists.[258] Thus, in the colloidal suspension of arsenious sulphide (prepared as described on p. 125) the ‹sulphide particles› carry ‹negative› charges, and the solution bathing them is ‹positive›. Conversely, colloidal ‹ferric hydroxide› (p. 127), in water, is charged with ‹positive›, the water with ‹negative›, electricity. The existence and character of these charges may be readily demonstrated by the passage of a current of electricity through the colloidal suspension in U tubes (‹exp.›).[259] Colloidal arsenious sulphide is found to migrate toward the positive pole, ferric hydroxide [p132] to the negative pole, the movement being easily followed by the color of the suspended particles.

The following colloids, which are of interest in analytical chemistry, are found to carry a ‹negative› charge in pure water: Colloidal ‹acids› (silicic, stannic), ‹sulphides› (As_{2}S_{3}, As_{2}S_{5}, CdS, etc.), ‹salts› (AgI, AgCl) and ‹metals› (Au, Pt, Ag). It is noteworthy that the suspensions of finely divided clay, kaolin, quartz, carbon, carry the same charge as these colloidal suspensions. On the other hand, ‹metal hydroxides› (ferric, aluminium, chromic), and ‹basic› substances in general, carry ‹positive› charges.

On the other hand, the colloids of one important group are found to be almost without any electric charges;[260] the passage of an electric current through their suspensions has little or no effect on them. Perfectly neutral albumen and gelatine are common representatives of this group.

The charge on a colloid seems to be liable to variation with the nature of the liquid in which it is suspended. Colloidal platinum in water is negative, in a mixture of water and alcohol, positive.[2] Of peculiar interest and importance is the fact that some colloids ‹change› the ‹character› of their ‹charge› when the liquid, in which they are suspended, is made to pass from an ‹acid› to an ‹alkaline› condition, and ‹vice versa›.[261] For instance, albumen[262], colloidal silicic acid[263] and colloidal stannic acid[264] are negative in alkaline liquids, positive in acid. The relation of these facts to the chemical nature of the substances will be discussed presently.

«The Source of the Electrical Charges on Colloids.»—Different views are held as to the source of electrification of colloidal suspensions. Only the two views of most direct interest in analysis can be considered here. In the first place, it is possible that ‹partial ionization› of the colloidal suspension produces the electrical charge. It is a significant fact that ‹basic colloids› (metal hydroxides, some basic dyes) receive a ‹positive› charge, such as would be left, if they were slightly ionized as bases (or salts of bases) and insoluble (suspended) [p133] ‹positive ions were retained by the particles›. That would be the case, for instance, if particles of colloidal ferric hydroxide, [Fe(OH)_{3}]_{‹x›}, sent a few hydroxide ions into solution and the suspended particles included positive insoluble ions. ‹Acid› colloids, on the other hand, assume a ‹negative› charge, as would be expected from this point of view. Then, the behavior of silicic acid is particularly suggestive; in alkaline or ‹slightly› acid liquids, in which its ‹ionization as a weak acid› is favored or predominant (Chapter X), it carries a ‹negative› charge, the charge that would be left on it, if it ionized, in part, as an acid or its salt. Silicic acid shows some slight tendency to ionize also as a base (see Chapter X) and the basic form of ionization would be favored by the presence of a strong acid (Chapter X). Colloidal silicic acid, as stated above, ‹changes its charge› from negative to positive as the solution passes from an alkaline to a (strong) acid reaction,[265] just as if, in the acid liquid, it ionized slightly as a base (salt of a base) and ‹retained a difficultly soluble positive ion›. Albumen, which has the property of being both a base and an acid,[266] shows the same change of charge[267] in the colloidal condition and the change has been ascribed[268] to the change of basic and acid functions.

According to the other of the two theories, here considered, the electrification of the colloid may result from what is known as contact or surface electricity.[269] At the surface of two different substances there is always a potential difference.[270] In the case of finely divided suspensions, like the colloids, the contact surfaces are enormous, as compared with the surfaces involved in ordinary contact. Whether metals (Au, Pt, Ag) owe their charges to simple contact effects or to their (minimal) tendency to ionize (Chapters XIV and XV) is not known. In fact, no exact knowledge of the source of electrification of any colloid has yet been obtained.

«Precipitation of Colloids by Electrolytes and by Colloids.»—Substances in the colloidal condition, ‹which carry an electrical charge›, are readily precipitated by the addition of electrolytes to the colloidal suspensions (see the behavior of arsenious sulphide, p. 126). Negatively charged colloids are precipitated by the action of positive ions, positively charged colloids by the action of negative ions (Hardy's rule[271]). The precipitated colloid carries [p134] with it a part of the precipitating ion[272] (‹adsorption›) and the weights of ions, carried down by a given quantity of a given colloid, are proportional to the equivalent weights of the ions.[273] The precipitation thus appears to be intimately associated with the neutralization of the charge on the colloid. In accordance with this conclusion, it has also been found that a colloid may be precipitated by a colloid carrying the opposite charge.[274] Thus, colloidal arsenious sulphide, carrying a negative charge, and colloidal ferric hydroxide, carrying a positive charge, mutually precipitate each other (‹exp.›[275]).

The purple precipitate (‹purple of Cassius›), which is formed when stannous chloride is added to gold chloride (p. 126), contains ‹colloidal gold›,[276] which, in suspension, is charged with negative electricity, and ‹colloidal stannic acid›,[276] which in acid solution, presumably, carries a positive charge[277]: these two colloids mutually precipitate each other in the presence of hydrochloric acid.[278]

When the precipitate is treated with an alkaline liquid (ammonium hydroxide solution), the charge on the stannic acid becomes negative, both colloids acquire the same charge and the precipitate dissolves, to form the beautifully tinted suspensions of colloidal gold (p. 126). In this condition the colloids (gold and stannic acid) are sensitive to precipitating electrolytes, and the solution is more sensitive to a mixture of magnesium nitrate and ammonium nitrate than to ammonium nitrate alone (‹exp.›), as is to be expected from a ‹negative› suspension (see below).

The characteristic difference in behavior between ‹ortho-›, ‹pyro-› and ‹metaphosphoric acids› toward albumen, which is used as a characteristic analytical test to distinguish metaphosphoric acid from the other two acids,[279] appears to be due to similar relations[280]: metaphosphoric acid, which precipitates (coagulates) albumen, is colloidal, ortho- and pyrophosphoric acids are not. The [p135] coagulation seems to be the result of the union of the ‹negative colloid› metaphosphoric acid (or a complex negative colloidal ion thereof) with the ‹positive colloid› albumen (or a positive colloidal ion thereof).

«The Precipitating Power of Electrolytes and the Valence of their Ions.»—The complete precipitation of colloids, which carry electric charges, depends on the concentrations of the colloid and the electrolyte; in this connection the important observation has been made that the ‹precipitating power› of electrolytes ‹increases decidedly with the valence of the precipitating ions› (H. Schulze's rule[281]). Bivalent ions are far more efficient than univalent; trivalent, in turn, still more effective than bivalent.

Thus, colloidal As_{2}S_{3}, carrying a negative charge, is precipitated by the positive ions of added electrolytes. The addition of a few drops of a molar solution of ammonium nitrate (the precipitating ion is NH_{4}^{+}) to 20 c.c. of the colloidal suspension[282] produces a slight precipitate; complete precipitation requires 3.5 to 3.6 c.c. of the ammonium nitrate solution.[283] Only 0.06 c.c. of an equivalent solution[284] of magnesium nitrate (the precipitating ion is Mg^{2+}) is required, and as little as 0.015 c.c. of an equivalent solution of aluminium nitrate[285] (the precipitating ion is Al^{3+}) has the same effect (‹exp.›). An increase in valence of the ‹negative› ion, which is not the precipitating ion in this case, does not affect the result appreciably: 3.5 c.c. of a solution[286] of ammonium sulphate, (NH_{4})_{2}SO_{4}, equivalent to the solution of NH_{4}NO_{3}, is also required for the complete precipitation of the colloidal As_{2}S_{3} (‹exp.›), although the one contains the univalent ion, NO_{3}^{−}, the other the bivalent ion, SO_{4}^{2−}.

Conversely, a ‹positively› charged colloid, like ferric hydroxide, may be precipitated by much smaller quantities[287] of bivalent negative ions than of univalent ions, etc. [p136]

«Nonprecipitation of Nonelectrified Colloids by Electrolytes.»—Colloids which do not carry any electric charges of moment (see p. 132) are also not precipitated by dilute solutions of electrolytes. Heat, the addition of concentrated salt solutions in great excess (whose effect is probably a dehydrating one), or of other solvents (‹e.g.› alcohol), are the agents most commonly used to effect coagulation in such cases.

«Protective Action of Colloids on Other Colloids.»—Colloids, particularly such as are not sensitive to precipitation by electrolytes, increase in a remarkable degree the stability of the colloidal condition of substances, that carry electrical charges and, ordinarily, would be very sensitive to precipitation. Thus, small quantities of albumen are used to render colloidal silver preparations more stable. Tannic acid, gelatine and albumen and related compounds are frequently used as such protective agents. It is supposed that they form protecting films around the colloidal particles.

«Applications in Analysis.»—From the preceding discussion we may now draw the following conclusions concerning the consideration that is to be given to the colloidal condition as a factor in qualitative analysis. The ‹absence of electrolytes› in solution must favor the production of the colloidal condition, which would result in the nonprecipitation or "solution" and ‹consequent loss of› substances, which it is intended to precipitate. Such absence of electrolytes in solution is most likely to be met with, in the first place, in the ‹washing› out of ‹precipitates›. When the larger part of the mother liquor is washed away by the use of pure water, many precipitates show a tendency to "run through a filter," forming colloidal suspensions in the pure water in the filter and being again precipitated as the colloid mixes with the electrolytes in the filtrate. Precipitates, showing this tendency to assume the colloidal condition, are therefore washed with appropriate solutions of electrolytes, rather than with pure water. Ammonium nitrate solution is most frequently available in qualitative analysis, because neither the ammonium-ion nor the nitrate-ion tends to interfere with the subsequent examination of the solution. When chloride-ion is not likely to interfere, ammonium chloride may be used. Thus, the sulphides of the arsenic, copper and zinc groups are washed with solutions containing ammonium nitrate (the [p137] chloride may be used for the zinc group[288]) rather than with pure water (or pure hydrogen sulphide water[288]). In quantitative analysis, aluminium hydroxide is also washed with ammonium nitrate solution, silver chloride with acidulated (nitric acid) water, lead sulphate with dilute sulphuric acid, etc.

In the second place, if precipitations are attempted either in rather dilute solutions or in solutions of little ionized substances (arsenious acid and hydrogen sulphide), the addition of an electrolyte is frequently required ‹to insure precipitation›. Thus, the presence of ammonium chloride, or nitrate, in excess, is helpful in the precipitation of the sulphides of the zinc group; the addition of hydrochloric acid (or other electrolyte) is required to effect the precipitation of arsenious sulphide from a solution of the oxide (p. 126).

In the next place, account must be taken, in analytical work, of the fact that ‹colloids carry down› with them the ‹precipitating ion› by which they are coagulated, a fact which may lead to the ‹loss of ions› which, it is intended, should be kept in solution. To a certain extent, this loss may also be avoided by insuring the presence of electrolytes (acids, ammonium salts) in sufficient concentration to cause the coagulation without the aid of the ions which, it is intended, should not be precipitated. In view of the much weaker precipitating power of univalent ions (of hydrochloric acid, ammonium nitrate and chloride), as compared with that of polyvalent ions, which may be present, the acid and ammonium salts must not be used in too small concentrations. In quantitative analysis, when conditions permit it, ammonium or sodium sulphate is frequently substituted for the ammonium salts of the univalent monobasic acids. The washing of the precipitated colloid with such salt solutions gradually removes[289] the ions which are precipitated with the colloid and forms a further safeguard against their loss. But this source of loss is avoided only with great difficulty and is seldom absolutely removed.

Finally, the presence of ‹protective colloids›, especially of the [p138] gelatine and albumen type, may interfere so decidedly with the common precipitation tests for ions, that ‹their destruction is imperative›, before these tests can be applied with any degree of confidence. Thus, the mixing of solutions (0.1 molar) of silver nitrate and hydrochloric acid, each containing one per cent of gelatine, fails to produce the ordinary, characteristic precipitate[290] of silver chloride, ‹the reaction which is used to determine the presence of the silver-ion in systematic analysis›.

The mixture is opalescent and, in reflected light, looks opaque-white; on somewhat prolonged standing a white ‹milk› is produced, but no precipitate. When the mixture is boiled, the same deep white milk is formed, but no coagulated precipitate, the mixture running unchanged through a filter. Hydrogen sulphide converts the mixture into a similar suspension of the black sulphide.

FOOTNOTES:

[227] The ratio is affected somewhat by the fineness of division of liquids and solids as a result of surface tension phenomena, as explained below.

[228] An aqueous solution of iodine and potassium iodide shaken with chloroform gives similar results, and the difference in color between the two layers is an advantage for a lecture experiment. But the iodine is partially combined with the iodide, according to KI + I_{2} ⇄ KI_{3}, or I^{−} + I_{2} ⇄ I_{3}^{−}, and the theoretical relations are not so simple as for bromine in aqueous and chloroform solutions.

[229] To accelerate the action, the mixture is shaken vigorously. After the separation of the layers, the bromine may be recognized in the aqueous layer by its color, or by the addition of potassium iodide (liberation of iodine).

[230] [Br]_{aq.} and [Br]_{ch.} are used to express the actual concentrations at any given moment we wish to consider, for instance at the moment equilibrium is reached. The ratio ‹k›_{2} / ‹k›_{1} for any substance S is found to be equal to the ratio of the ‹solubilities› of the substance in the two solvents. That it must be so can be proved by applying the law of physical equilibrium to the mixed solvents in contact with an excess of the substance, ‹i.e.› to its ‹saturated› solutions (see below).

[231] Similar relations hold for the condition of equilibrium between a liquid and its vapor.

[232] Curie, ‹Bull. Soc. Min.›, «8», 145 (1885); Ostwald, ‹Grundlagen der Anal. Chem.›, p. 22 (1894).

[233] Hulett, ‹Z. Phys. Chem.›, «37», 384 (1901). See also Ostwald, ‹ibid.›, «34», 495 (1900), on the solubilities of finely divided mercuric oxide ("yellow oxide") and of larger crystals ("red oxide").

[234] A 1 / 1000 solution of AuCl_{3}, 2 aq., and a 1.2 / 1000 solution of SnCl_{2}, 2 aq., may be used conveniently. When equal volumes of the solutions are mixed, the desired precipitate is formed.

[235] Stannous chloride is usually sufficiently contaminated with stannic salt. Add a few drops of bromine- or chlorine-water to 100 c.c. of a ‹pure› stannous chloride solution.

[236] The action is an exceedingly sensitive qualitative ‹test for gold›. By a modification of the test Donau was able to detect as little as 2E−9 gram of gold (‹Monatshefte f. Chem.›, «25», 545 (1904)).

[237] The nature of the reduction reaction is discussed in Chapters XIV and XV.

[238] Faraday, ‹Proc. Royal Soc.›, «8», 356 (1857), ‹Phil. Mag.› (4), «13», 401 (1857) (Stud.).

[239] Zsigmondy, ‹Liebig's Annalen›, «301», 30 (1898).

[240] Bredig, ‹Z. f. Elektrochem.›, «4», 514, 547 (1898), and ‹Anorganische Fermente›, 1901. Colloidal preparations of platinum, silver and many other metals may be prepared in the same way. In ether, colloidal preparations of the alkali metals may be made (Svedberg, ‹Ber. d. chem. Ges.›, «38», 3616 (1905), «39», 1708 (1906)).

[241] ‹Argentum Credé› may be conveniently used. It contains, with the metallic silver, a small percentage of albumen, which is added for reasons discussed below. Brown solutions are formed at once.

[242] First a thin ‹film of› black ‹carbon› is produced ‹round the metal›, then the latter appears in the form of a filigree of silver.

[243] For details of the preparation, see A. A. Noyes, ‹J. Amer. Chem. Soc.›, «27», 94 (1905).

[244] This process of separation of substances, which do not pass through membranes, from such as do, is called ‹dialysis›. It was first used by Graham, ‹Trans. Royal Soc.›, London, «151», 183–224 (1861) («Stud.»).

[245] Graham made the first extended investigations in this field: ‹Trans. Royal Soc.›, «151», 183 (1861); ‹J. Chem. Soc.› (London), «17», 318 (1864) («Stud.»). He found that amorphous, gelatinous bodies like ferric hydroxide, aluminium hydroxide, silicic acid, gelatine, glue, dextrin, caramel, albumen and similar bodies do not pass through membranes and may be obtained by dialysis in the ‹colloidal› condition. Such substances were called "colloids" by Graham, the name referring to the Latin for gelatine. Substances which pass through membranes readily were found by Graham to resemble in behavior such bodies as are crystallizable when solid; such compounds were classified by him as "crystalloids." That liquids containing substances in the colloidal condition (‹e.g.› arsenious sulphide, gold, silver and many other substances) may be prepared by methods other than dialysis, was found later by many investigators and, in a few cases, previous to Graham, ‹e.g.› by Faraday, ‹loc. cit.› A brief history of the chemistry of colloids is found as an introduction to Wo. Ostwald's ‹Kolloidchemie›, pp. 1–63 (1909).

[246] ‹Cf.› Wo. Ostwald, ‹loc. cit.›, p. 193.

[247] Before Graham's time, and for the few colloidal liquids then known, this view was held by such men as J. B. Richter, Berzelius and Faraday (‹loc. cit.›) (‹cf.› Wo. Ostwald, ‹loc. cit.›, p. 19). The first extended experimental investigation in support of it was made by Barus and Schneider, ‹Z. phys. Chem.›, «8», 278 (1891). Bredig was also an early and consistent champion of this view (‹vide› his ‹Anorganische Fermente›, 1901).

[248] ‹Cf.› Wo. Ostwald, ‹loc. cit.›, pp. 102–114. Graham considered "colloidal silicic acid a liquid miscible with water in all proportions." According to modern ideas, no true miscibility exists, but a suspension or emulsion is formed (see Ostwald, p. 237).

[249] Siedentopf and Zsigmondy, ‹Drude's Annalen›, «10», 1 (1903). Zsigmondy, ‹Colloids and the Ultramicroscope› (1909), Chapter V.

[250] Zsigmondy, ‹Z. für Elektrochem.›, «8», 684 (1902); Siedentopf and Zsigmondy, ‹loc. cit.›

[251] Zsigmondy, ‹loc. cit.›, p. 161. A µµ is 1E−6 mm. The hydrogen molecule is considered to have a diameter of 0.1 µµ (O. E. Meyer), the alcohol molecule one of 0.5 µµ (Zsigmondy, ‹loc. cit.›, plate IV, p. 157).

[252] Weimarn, ‹Chem. Zentralblatt›, 1907, II, p. 1293.

[253] Paal, ‹Ber. d. chem. Ges.›, «39», 1436, 2859 (1906).

[254] Other varieties of heterogeneous colloidal mixtures are tabulated by Wo. Ostwald, ‹loc. cit.›, p. 96.

[255] The "Cassius' purple" test for gold is an instance where the colloidal condition is used in analysis for a positive test. See Wo. Ostwald, ‹loc. cit.›, p. 68, for other, similar applications for positive tests.

[256] ‹Vide› Wo. Ostwald's ‹Kolloidchemie›, 1909, and the references given by Noyes, ‹loc. cit.›, p. 86.

[257] A general discussion of the preparation and properties of colloidal mixtures is given by A. A. Noyes, ‹J. Am. Chem. Soc.›, «27», 86–104 (1905) («Stud.»).

[258] Picton and Linder, ‹J. Chem. Soc.› (London), «61», 160 (1892), «67», 63 (1895), «71», 568 (1897), etc., and others. Wo. Ostwald, ‹loc. cit.›, p. 240, gives a list of references.

[259] The arrangement of the experiment is described by Noyes, ‹loc. cit.›, p. 98.

[260] ‹Cf.› Wo. Ostwald, ‹loc. cit.›, p. 108. Billitzer has found that gelatine is positive in acid solution, negative in alkaline, ‹Z. phys. Chem.›, «51», 147 (1905). The charges are, however, relatively small ones.

[261] Billitzer, ‹Z. f. Elektrochem.›, «8», 638 (1902). This is probably true of all amphoteric colloids (Chapter X); it is also true of many other substances, which are not pronouncedly amphoteric. (‹Cf.› Perrin, ‹Comp. rend.›, «136», 1388 (1903); Billitzer, ‹Z. phys. Chem.›, «51», 157 (1905).)

[262] Hardy, ‹J. of Physiology›, 24, 288 (1899); ‹Z. phys. Chem.›, «33», 387 (1900).

[263] Billitzer, ‹loc. cit.›, p. 159; Müller's ‹Allgemeine Chemie der Kolloide›, 1907, p. 79.

[264] See below, p. 134.

[265] In a slightly acid solution colloidal silicic acid is negatively charged; in a ‹strong› acid solution, positively—a relation which agrees with its ‹predominantly acid character›.

[266] The general class of substances, showing both basic and acid properties, of which albumen is a derivative, is described in a footnote on glycocoll, Chapter X, p. 188.

[267] Hardy, ‹loc. cit.›

[268] J. Loeb, University of California Publications, ‹Physiology›, «2», 149 (1904).

[269] Very little is known about the nature of contact electricity. It is even doubtful whether it is different, in principle, from ionization.

[270] W. Ostwald, Lehrbuch der Chem., «2», (1) 553 (1903).

[271] Hardy, ‹Proc. Royal Soc.›, «66», 110 (1899); ‹Z. phys. Chem.›, «33», 391 (1900).

[272] Picton and Linder, ‹J. Chem. Soc.› (London), «67», 63 (1895).

[273] Whitney and Ober, ‹J. Am. Chem. Soc.›, «23», 852–856 (1901) (Stud.).

[274] Picton and Linder, ‹loc. cit.› «71», 572 (1897); Lottermoser, ‹Anorganische Kolloide›, p. 76; Biltz, ‹Ber. d. chem. Ges.›, «37», 1095 (1904).

[275] Precipitation is complete only when the colloids are used in the proportions required to neutralize each other's charges [Billitzer, ‹Z. phys. Chem›., «51», 140 (1905)]. The proportions to be used must be determined in each case, most simply by trial (Noyes, ‹loc. cit.›, p. 101), but quantitative methods for determining the charges, by titration, are also known (‹cf.› Billitzer, ‹loc. cit.›).

[276] E. A. Schneider, ‹Z. anorg. Chem.›, «5», 80 (1894).

[277] See the above discussion on silicic acid. Stannic acid has a greater tendency to form a base than has silicic acid.

[278] Zsigmondy, ‹Liebig's Annalen›, «301», 361 (1898).

[279] ‹Cf.› Fresenius, p. 334, or Smith's ‹Inorganic Chemistry›, p. 468.

[280] Mylius, ‹Ber. d. chem. Ges.›, «36», 775 (1903); Biltz, ‹ibid.›, «37», 1116 (1904).

[281] ‹J. für prakt. Chem.›, «25», 431 (1882).

[282] The suspension used is prepared by saturating, with hydrogen sulphide, an aqueous solution of arsenious oxide. The latter is saturated on a steam bath, cooled to 20°, filtered and diluted with an equal volume of water before it is used.

[283] Eight grams of NH_{4}NO_{3} per 100 c.c.

[284] 0.6 c.c. of a ‹tenth›-normal solution is used, containing 1.28 gram Mg(NO_{3})_{2}, 6 aq., in 100 c.c. Precipitation was found to be incomplete with 0.5 c.c.

[285] 0.15 c.c. of a ‹tenth›-normal solution is used, containing 1.2 gram Al(NO_{3})_{3}, 4 aq., in 100 c.c. Precipitation was found to be incomplete with 0.1 c.c. of the solution.

[286] 3.3 grams (NH_{4})_{2}SO_{4} in 100 c.c.

[287] Freundlich found, for instance, that NaCl, KCl, BaCl_{2}, in equivalent concentrations, had practically the same effect on colloidal ferric hydroxide, but only ‹one-fortieth› as much of a ‹sulphate› (the precipitating ion is SO_{4}^{2−} versus Cl^{−}) was required; ‹Z. phys. Chem.›, «44», 129 (1903).

[288] For other reasons (‹e.g.› to prevent oxidation of the sulphides), hydrogen sulphide is also used in the solution for washing the arsenic group, and ammonium sulphide in that for the zinc group (see Lab. Manual, pp. 101 and 110).

[289] Picton and Linder, ‹loc. cit.›, and Whitney and Ober, ‹loc. cit.›

[290] A. A. Noyes describes a similar experiment with sodium chloride and silver nitrate, ‹loc. cit.›

[p139]