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Ferric oxide, Fe2O3

Ferric oxide, Iron sesqui-oxide, Fe2O3, occurs in abundant quantities in nature, both in the massive and crystalline forms, the former being known as red hcematites, whilst the latter are termed specular iron when the crystals are rhombohedra and scalenohedra, or micaceous iron when in thin translucent scales.

In the laboratory ferric oxide may be obtained in a variety of ways. Thus when ferric hydroxide or sulphate is strongly heated ferric oxide remains behind, and the same applies if ferric chloride or sulphide, ferrous oxide or carbonate, or indeed the majority of ferrous salts, are heated in contact with air. Several of these methods are adopted on a manufacturing scale. For example, in the manufacture of sulphuric acid iron pyrites is roasted in air, leaving a residue of ferric oxide. Thus: -

2FeS2 + 11O = 4SO2 + Fe2O3.

Ferric oxide is manufactured in large quantities for use as a pigment by roasting ferrous sulphate obtained by weathering iron pyrites.

Ferric oxide is also prepared from liquors containing ferric salts in solution, and which are otherwise waste products in many manufacturing processes. If ferrous salts are present, they are first oxidised by addition of nitric acid or bleaching powder. The acid is neutralised by addition of soda or lime, ferric hydroxide being precipitated. The washed product is finally dehydrated by heat, becoming perfectly anhydrous at 500° C. upwards.

Ferric oxide, in a more or less impure condition, is manufactured for pigmentary purposes by the ignition of natural ferric hydroxides such as ochres. The colour of the final product depends largely upon the completeness with which the water has been expelled. Thus Venetian reds are produced after ignition at dull red heat for some eight hours; ten hours' heating yields light reds; twelve hours the so-called Indian reds; and so on.

Crystalline Ferric Oxide

Small crystals of ferric oxide may be obtained by fusing amorphous ferric oxide with sodium borate (borax), and extracting them from the cooled mass with the aid of dilute aqueous nitric or hydrochloric acid. Passage of gaseous hydrogen chloride over the ferric oxide at red heat; of steam containing some ammonium fluoride over the oxide at 600° C.; of ammonium chloride over ferric oxide at 700° C.; or of vaporised ferric chloride over heated lime, results in the formation of small crystals of ferric oxide. Crystals have been found, produced by the first of these methods, in iron pipes which have for many years conducted alternately hydrogen chloride and air in a plant connected with Deacon's process for the manufacture of chlorine. Crystals have also been found as products of smelting operations, whilst crystals having the form of specular iron have been found in iron rust from a building seven hundred to eight hundred years old. By heating mixed solutions of the sulphates of copper and iron (ferrous) in sealed tubes to 210° C. for ten hours, crystals resembling those of micaceous iron are produced. When ferric sulphate is heated, it decomposes, yielding hexagonal lamellae of ferric oxide, having a density of 4.95 at 14° C.

Ferric oxide, pseudomorphous with magnetite, may be obtained by heating magnetite crystals in the blow-pipe flame for several hours; oxidation takes place, the crystals retaining their original form almost unaltered, but losing their magnetic properties.

Ferric oxide may be obtained in an exceptionally pure condition, very useful for analytical purposes, by dissolving a piece of metallic iron, preferably of high purity, in hydrochloric acid, diluting, and precipitating as sulphide by passage of hydrogen sulphide, any copper, etc., with which the metal is contaminated. The clear filtrate is evaporated to small bulk, oxidised with nitric acid, evaporated to dryness with hydrochloric acid, taken up with water, and extracted with ether. The ethereal solution is distilled, and the residual ferrous and ferric chlorides dissolved in diluted hydrochloric acid, reduced with sulphur dioxide and precipitated as ferrous oxalate with the ammonium salt. After thorough washing the oxalate is ignited to ferric oxide.

Pure ferric oxide is recommended for the standardising of permanganate solutions for volumetric analysis. It is dissolved in hydrochloric acid, reduced with stannous chloride, and titrated with permanganate.

As obtained by any of the foregoing methods, ferric oxide is an extremely stable substance, soluble in acids only with difficulty. It melts at 1565° C., and freezes at 1562° to 1565° C.

Its cubical coefficient of expansion with rise of temperature is 0.00004.

Its specific heats for various temperature intervals are as follow: -

Temperature Interval. °C.Specific Heat.Molecular Heat.
-191.9 to -81.60.072811.59
-191.9 to -80.40.0724
-73.6 to 00.131821.05
-73.8 to 00.1318
3.9 to 43.80.160425.53
3.2 to 44.30.1596

At about 2500° to 3000° C. it shows signs of volatilising.

It crystallises in tabular hexagonal scales belonging to the hexagonal system, and possesses a steel-like lustre. The edges are ruby-red in colour, and give a red streak. Density, 5.187 to 5.193.

At 640° C. ferric oxide appears to undergo a polymorphic change. The magnetic properties of ferric oxide are ordinarily very feeble, but when heated to a very high temperature, as, for example, in the electric furnace, or oxy-hydrogen flame, it becomes magnetic, owing to conversion into the magnetic oxide, Fe3O4. It is not reducible by solid carbon, in the absence of gases, below 950° C. When heated with platinum to 1600° C. in the air, ferric oxide is reduced to the metal, oxygen being evolved and the iron passing into solid solution in the platinum. Under low oxygen pressures the temperature of reduction may be as low as 1200° C.

When heated to 480° C., under a pressure of twelve atmospheres of oxygen, ferric oxide undergoes no chemical change, a higher oxide not being formed.

With magnetic oxide ferric oxide yields a continuous series of solid solutions, ranging from Fe2O3 down to practically Fe3O4 itself. No intermediate oxides have been detected.

The dissociation pressures of ferric oxide between 1100° and 1400° C. are as follow: -

Temperature.°C.Pressure (mm. mercury).

Ferric oxide is reduced to ferrous oxide or the metal, according to circumstances, by carbon monoxide. At temperatures below 1000° C. the dry gas is more effective than the moist, but at 1050° C. both moist and dry gases behave alike. At and below 850° C. the iron is converted into carbide. At 700° C. the reaction

3Fe2O3 + CO = 2Fe3O4 + CO2

also takes place.

Ferric oxide is reduced to metallic iron when strongly heated in a current of hydrogen, the water vapour formed during the reduction being rapidly carried away in the current of gas.

The temperature at which reduction begins depends on the temperature to which the oxide has previously been heated.

In the case of hydrogen gas the reaction usually begins at about 330° C. with the formation of magnetic oxide. Thus: -

3Fe2O3 + H2 = 2Fe3O4 + H2O.

Carbon monoxide is active even at 240° C. At 500° C. ferrous oxide is formed: -

Fe3O4 + H2 = 3FeO + H2O;

and at 600° C. complete reduction to metallic iron is effected: -

FeO + H2 = Fe + H2O.

The metal obtained at this temperature, however, is not pyrophoric.

It must be remembered that the foregoing temperatures are only approximations, and refer to experiments of ordinary duration. For example, the last reaction, stated to begin at about 600° C., will take place, albeit very slowly, at considerably lower temperatures. Thus, after ninety-six hours of treatment, Moissan was able to reduce the oxide to metallic iron at 440° C., and the metal was then pyrophoric.

If, however, the reduction is made to take place in a more or less confined area, various equilibria are set up, according to circumstances. Thus

Fe2O3 + 3H2 ⇔ 2Fe + 3H2O, Fe2O3 + 3CO ⇔ 2Fe + 3CO2,

and these reactions are still further complicated by the formation of ferrous oxide and magnetic oxide.

Ferric oxide is rapidly reduced by nascent hydrogen. If added to hydrochloric acid in which metallic iron is dissolving, it is reduced and quickly dissolved, yielding ferrous chloride.

Ferric oxide is not attacked by thionyl chloride, SOCl2, at the ordinary temperature. At 150° C. the following reaction readily takes place: -

Fe2O3 + 3SOCl2 = 2FeCl3 + 3SO2,

the ferric chloride crystallising out in green hexagonal plates.

Heated with sulphur, ferric oxide yields ferrous sulphide and sulphur dioxide. With hydrogen sulphide at white heat hydrogen and sulphur dioxide are evolved: -

2Fe2O3 + 7H2S = 4FeS + 3SO2 + 7H2.

Ferric oxide reacts slowly at 700° to 800° C. with sulphur dioxide, yielding sulphur trioxide and magnetic oxide: -

3Fe2O3 + SO2 = 2Fe3O4 + SO3.

Below 600° C. there is no action.

Ferric oxide is attacked by chlorine at high temperatures, yielding a sublimate of ferric chloride, the same salt being also produced upon ignition of the oxide in hydrogen chloride. Ammonia, under similar conditions, is oxidised to water, a nitride of iron being produced.

Ferric oxide is dissolved by hydrochloric acid. Nitric acid does not attack the ignited oxide. Sulphuric acid, particularly a mixture of 8 parts of acid with 3 parts of water, effects its solution.

When heated with calcium sulphate, ferric oxide causes the expulsion of sulphur trioxide, and a similar reaction takes place with lead and magnesium sulphates.

The heats of formation of anhydrous ferric oxide below 400° C. are given as follow: -

2[Fe] + 3(O) = [Fe2O3] + 194,400 calories.
2[Fe] + 3(O) = [Fe2O8] + 192,200 calories.
2[FeO] + (O) = [Fe2O3] + 65,200 calories.
2[FeO] + (O) = [Fe2O3] + 63,700 calories.
2[Fe3O4] + (O) = 3[Fe2O3] + 54,500 calories.

Ferric Oxide as a Catalyst

Ferric oxide possesses the power of catalytically promoting the combination of sulphur dioxide and oxygen at red heat. The action is perceptible at temperatures just above 400° C., attaining a maximum at 625° C. when 70 per cent, of the sulphur dioxide is converted into trioxide. The origin of the ferric oxide is of considerable importance, that prepared from the hydroxide being particularly active. Admixture of copper oxide increases the efficiency, as does also the presence of arsenic at temperatures above 700° C.

In commercial practice, the ferric oxide is obtained from pyrites cinder, and experience shows that its catalytic activity is greatest if the ferric oxide is used fresh and not first allowed to get cold. It appears that if the air required for oxidation of the pyrites is dried previous to admission to the burners, the resulting mixture of sulphur dioxide and air is more sensitive to the action of the catalyst than is otherwise the case. Under highly favourable conditions 90 per cent, of the sulphur dioxide is converted into sulphur trioxide, and a conversion of 60 per cent, is quite readily obtained.

Owing to the high temperature required, however, there appears to be a loss of sulphur trioxide through dissociation when the process is carried out on a commercial scale. A process has therefore been patented according to which the conversion of the sulphur dioxide into trioxide is rendered practically complete by passing the partially converted gases emerging from the ferric oxide chamber through a second chamber containing platinum as catalyst.

The manner in which the ferric oxide is able to function as a catalyst has been the subject of discussion. According to one theory, it is supposed to undergo alternate reduction to magnetic oxide and oxidation to ferric oxide; thus: -

3Fe2O3 + SO2 = 2Fe3O4 + SO3,
4Fe3O4 + O2 = 6Fe2O3.

It has also been suggested that sulphur dioxide and oxygen combine with the ferric oxide to form ferric sulphate, which then dissociates into the trioxide, regenerating the ferric oxide. Thus: -

2Fe2O3 + 6SO2 + 3O2 = 2Fe2(SO4)3,
Fe2(SO4)3 = Fe2O3 + 3SO3.

A third theory has been put forward by Keppeler, according to which the catalytic action is attributable to a physical cause, the gases being condensed on the surface of the ferric oxide and uniting under these conditions. This theory explains the similarity existing between the catalytic activities of ferric oxide and platinum, and harmonises very satisfactorily with the conclusions of Bone and Wheeler, relative to the catalytic activity of ferric oxide and other heated surfaces in the surface combustion of mixtures of hydrogen and oxygen.

Many other reactions are known in wrhich ferric oxide acts as a catalyser. Thus, it assists the decomposition of mercuric oxide at temperatures between 360° and 480° C., and the oxidation of carbon monoxide with steam, the reaction proceeding very slowly at 250° C. but rapidly at 400° C.

It assists catalytically the conversion of sodium chloride into sodium sulphate when heated in a current of air along with finely divided pyrites. It also accelerates the decomposition of potassium chlorate when heated with this salt, much in the same way as manganese dioxide has long been known to do. More chlorine is evolved during the decomposition, however, and under certain conditions the oxygen is evolved at an even lower temperature than when manganese dioxide is used.

Polymorphism of Ferric oxide

The yellow colour in certain bricks is stated to be due to a yellow modification of anhydrous ferric oxide rendered stable by alumina.

It may well be, however, that the colour of ferric oxide is determined by the size of the grain rather than by any variation in molecular structure. Thus, brown and violet samples of ferric oxide are converted into the yellowish red variety by alternate grinding and washing - processes which, in view of the chemical stability of ferric oxide, are hardly likely to effect a molecular transformation.

Hydrated ferric oxide, Fe(OH)3

Hydrated ferric oxide, Fe(OH)3.Aq., occurs in nature in various stages of hydration, the best-known minerals being as follow: -

Turgite, 2Fe2O3.H2O.
Goethite, Fe2O3.H2O.
Hydrogoethite, 3Fe2O3.4H2O.
Limonite, 2Fe2O3.3H2O.
Xanthosiderite, Fe2O3.2H2O.
Limnite, Fe2O3.3H2O.
Esmeraldaite, Fe2O3.4H2O.

When heated with water under enormous pressures (some 5000 atmospheres) ferric oxide becomes hydrated. At 42.5° C. it yields brown iron stone; at 42.5° to 62.5° C., Goethite, Fe2O3.H2O, and at higher temperatures, a hydro-haematite, 2Fe2O3.H2O, resembling turgite in composition.

Ferric hydroxide, dried at 100° C., gradually becomes rehydrated upon prolonged exposure to a saturated atmosphere. Although the above substances are usually described as hydrates of ferric oxide, it is by no means certain that all of them are to be regarded as definite chemical entities. They are mostly hygroscopic substances, the amount of water they contain at any moment fluctuating with the temperature and humidity of the atmosphere. The task of determining precisely how much of the contained water is merely physically attached to the oxide, and how much is chemically combined with it, is not easy.

Precipitated ferric hydroxide becomes gradually dehydrated when its temperature is raised. At 55° C. it attains the composition Fe2O3.3H2O, and with further rise in temperature, more water is gradually evolved until at about 385° C., a substance represented by the formula 10Fe2O3.H2O is obtained, which remains constant in weight for several hours at 385° to 415° C. At 500° C. the oxide is perfectly dehydrated, and its weight remains constant.

The subhydrate, 2Fe2O3.H2O, occurs in nature as the mineral turgite. It may be prepared in the laboratory by boiling ferric hydroxide with distilled water for prolonged periods, when it becomes partially dehydrated and loses its gelatinous appearance. It still contains from 4 to 6 per cent, of water, approximating to a composition of 2Fe2O3.H2O. The same effect is produced by heating the hydroxide for 1000 to 2000 hours at 50° to 60° C., the resulting oxide having a brick-red appearance, and a density of 4.5, and generally resembling red haematite. It seems reasonable to suppose that the change may take place at still lower temperatures, given sufficient time, and it is thus unnecessary to postulate the need for high temperatures to account for the production of some of our deposits of red haematite.

The monohydrate, Fe2O3.H2O, occurs in nature as the mineral Goethite (see p. 18), which is regarded as a definite crystalline hydroxide and not as a colloid like limonite. The monohydrate may be prepared in the laboratory in the amorphous condition by the prolonged boiling of the brown precipitate resulting from the addition of alkali to ferric chloride solution. The colour gradually changes from brown to brick- red, and the resulting hydrate is remarkably resistant to acid attack. Boiling concentrated nitric acid has but little effect, and even concentrated hydrochloric acid only attacks it after prolonged digestion at the boiling-point. After some hours of treatment with acetic acid at 100° C. a colourless colloidal solution is obtained, from which addition of a trace of sulphuric acid effects the precipitation of the insoluble monohydrate. If, however, the solution is prepared in the cold, it possesses a wine-red colour and reacts like a ferric salt.

When iron is exposed to moist air, it readily becomes coated with a light, friable, and porous mass of oxide, which is brown in colour, and is generally termed rust. Numerous analyses of rust have been published from time to time, which prove that its composition varies according to the age and method of formation of the rust. When formed by exposure to ordinary air its composition corresponds fairly closely to the formula Fe2O3.H2O.

The monohydrate has been obtained in crystalline form in several ways, such as, for example, by the action of concentrated potassium hydroxide solution upon potassium nitroprusside.

When iron is attacked by fused sodium peroxide, dark red, tabular crystals of a monohydrate, Fe2O3.H2O, are obtained, of density 3.8 at 27° C. The hydrate is magnetic, and when heated to low redness a magnetic form of anhydrous ferric oxide is obtained.

Crystals of Goethite have been maintained under a pressure of 9500 atmospheres for 26 days at 15° C. without showing any loss of water, although almost all the water is expelled when the crystals are immersed in water for 7 days at 320° to 330° C. under a pressure of 135 atmospheres.

An unstable variety of Fe2O3.H2O is obtained by partially dehydrating the dihydrate by exposure in a desiccator.

Ferrous acid, Fe2O3.H2O, or HFeO2, possesses the same empirical formula as the monohydrate.

The sesqui-hydrate, 2Fe2O3.3H2O, is generally regarded as resulting when the normal trihydrate is allowed to dry in a vacuum. It occurs in nature as the mineral limonite, and is a valuable source of iron. It is probably not a chemical entity but a colloidal substance containing adsorbed water.

A satisfactory explanation for the yellow colour of limonite has not as yet been offered. Possibly the colour is due to the adsorption of an iron salt, but this has not been proved.

A substance having a composition corresponding to the dihydrate, Fe2O3.2H2O, is obtained by dissolving reduced iron in hot, diluted sulphuric acid and heating until the acid begins to fume and the iron is transformed into a faintly red, crystalline powder. The acid is poured off and the precipitate shaken with sodium hydroxide solution and finally washed with water. The substance is formed as brownish plates of density 3.234 at 15° C.

It is probably not a definite hydrate. It loses water when kept in a desiccator, yielding a substance corresponding in composition to the monohydrate, Fe2O3.H2O, but which is not very stable, and is probably not the normal monohydrate.

The trihydrate, Fe2O3.3H2O, or Fe(OH)3, is the brown voluminous precipitate obtained on adding an alkali to a solution of a ferric salt. The reaction is delicate, 1 part of iron in 80,000 parts of water being detectable by the precipitate obtained if ammonia is the alkali employed.

A mixed solution of ferrous sulphate and hydrogen carbonate deposits ferric hydroxide upon standing, even in the dark. Its heat of formation is

2[Fe] + 3(O) + 3H2O = 2[Fe(OH)3] + 191,150 calories.

If the trihydrate is boiled with water for several hours it gradually assumes a brick-red colour, being converted into the monohydrate, Fe2O3.H2O. The trihydrate is readily soluble in acids yielding ferric salts. It also is slightly soluble in ordinary distilled water, namely, to the extent of 0.151 mgm. of Fe(OH)3 per litre at 20° C.

Freshly precipitated ferric hydroxide adsorbs arsenious acid from solution; the extent of adsorption is diminished by the presence of sodium hydroxide, but addition of sodium chloride appears to have no influence. It has therefore been recommended as an antidote in cases of arsenic poisoning. The amount of adsorption is given by the expression

when у and x are the amounts of acid adsorbed and remaining in solution respectively. The following data illustrate the general agreement between the calculated and observed values of x: -

yx (calculated)x (observed)

Ferric hydroxide dissolves in fused sodium hydroxide, and, if the latter is present in excess, lustrous needles and lamellae of composition approximating to 20Fe2O3.32H2O.3Na2O are obtained. These dissolve readily in acids, and begin to lose water at 120° C., becoming anhydrous at dull red heat. The resulting product, 20Fe2O3.3Na2O, resembles specular haematite in appearance and is conceivably a solution of sodium ferrate in ferric oxide. By heating ferric hydroxide with sodium hydroxide to 110° C., a product resembling Goethite, Fe2O3.H2O, is obtained, but again admixed with small quantities of sodium oxide. Ferric hydroxide dissolves in potassium hydroxide in the presence of ozone, yielding potassium perferrate.

On passing a current of air into a hot, concentrated solution of sodium hydroxide containing ferric hydroxide in suspension, an appreciable quantity of iron passes into solution without colouring the liquid. On standing for several days the liquid becomes turbid owing to separation of ferric hydroxide, but is readily clarified again by warming. Probably sodium ferrate is formed in solution.

Ferric hydroxide is readily soluble in mineral acids yielding the corresponding ferric salt, and in ferric chloride solution yielding oxy compounds which are acidic in their behaviour, decomposing carbonates, and can therefore hardly be termed basic chlorides. It dissolves in aqueous oxalic acid to an extent directly proportional to the concentration of the acid, no definite basic oxalate being formed at 25° C. from solution. It does not combine with carbon dioxide when freshly precipitated and suspended in water.

It dissolves to a considerable extent in a concentrated solution of aluminium sulphate, yielding a brown solution which may be evaporated to dryness without decomposition. Addition of water induces the formation of a basic salt, 3Fe2O3.SO3.3H2O.

Inasmuch as the majority of ferric salts in the crystalline state are whitish, the question arises as to why ferric hydroxide should be brown. Ferric hydroxide can be obtained as a white precipitate on adding a freshly prepared and concentrated solution of a ferric salt to cooled ammonium hydroxide solution, but it rapidly becomes brown - due, it has been suggested, to molecular condensation or aggregation. When kept under water for a year, ferric hydroxide changes in colour from brown to yellowish red, and about 30 per cent, becomes insoluble in dilute acids. Possibly these changes are due to a similar cause.

A mineral having a composition corresponding to a tetrahydrate, Fe2O3.4H2O, has been found as an inclusion in limonite in Esmeralda Country, Nevada, whence the name Esmeraldaite. It is glassy and brittle, but possesses a yellowish brown streak. Its density is 2.58. A substance of similar composition is obtained when the voluminous precipitate of hydrated oxide, resulting from addition of ammonia to dilute ferric chloride solution, is allowed to dry in the air. It is vitreous in appearance, black in colour when viewed in mass, although thin sections appear red by transmitted light. Its density is 2.436 at 15° C. Pressure does not decompose it, but it loses water when placed in a desiccator.

Colloidal ferric hydroxide

Ferric hydroxide may be obtained in colloidal solution by adding 5 c.c, of 33 per cent, ferric chloride to a litre of boiling water and removing the chloride remaining, together with the hydrochloric acid, by dialysis.

Another method consists in boiling a solution of ferric nitrate with copper filings or zinc dust. The ferric nitrate need not be specially isolated for the purpose, but may be made merely as an intermediate product during the course of the reaction - if, for example, iron filings containing copper are treated with concentrated nitric acid. After dilution and filtration, the solution is dialysed, whereby a deep red liquid is obtained, containing colloidal ferric hydroxide.

By saturating a solution of ferric chloride with ammonium carbonate and purifying the solution thus obtained by dialysis, ferric hydroxide is readily obtained in the colloidal state, or by dialysing the solution obtained by adding ammonia to a solution of ferric chloride in such small quantities at a time that the ferric hydroxide at first thrown out is completely dissolved on stirring. The clear, dark brown solution so obtained scarcely tastes of iron. The last traces of chlorine are not removed in this way, although the colloid is free from ammonium salts.

When a 10 per cent, solution of ferric chloride is poured into excess of ammonia, the colloidal ferric hydroxide initially produced is coagulated by the ammonium chloride. On evaporating to dryness and washing with water, the ammonium salt washes out, and then the ferric hydroxide deflocculates, passing into subsequent wash waters as a red colloidal solution.

On addition of a dilute solution of potassium permanganate to one of ferrous chloride, and subjecting to dialysis, the pure colloidal hydroxide is readily obtained. The sol is also obtained by oxidising a solution of ferrous chloride containing one gram equivalent of FeCl2 per litre with a 3 per cent, solution of hydrogen peroxide.

The colloid, as usually prepared, is electro-positive in character, and may be precipitated from solution by electrolysis, by the addition of small quantities of electrolytes, or by the action of an oppositely charged colloid, such, for example, as (negative) arsenious sulphide, whereby the two electrical charges neutralise each other. The smallest quantities of a few electrolytes required to precipitate colloidal ferric hydroxide from solution are given in the following table: -

Electrolyte.Concentration in Gram-molecules per Litre of Solution.

From the above it is clear that it is the negative ion which influences the precipitation most, the divalent ions being considerably more effective than the monovalent.

Non-electrolytes, even in concentrated solution, have usually no action on the colloid.

When the colloidal solution is boiled with Fehling's solution (made by mixing copper sulphate and alkaline sodium potassium tartrate solutions), the colloid is precipitated along with cuprous oxide.

It is possible also to prepare colloidal ferric hydroxide with a negative charge. This may be done by adding slowly 100 c.c. of 0.01-normal ferric chloride solution to 150 c.c. of 0.01-normal sodium hydroxide, the mixture being continuously shaken during the process.

Ferric hydroxide is not unique in this respect, for stannic hydroxide has likewise been prepared both as a positive and a negative colloid.

Negative colloidal ferric hydroxide may be converted into the positive colloid by adding it to a very dilute solution of sodium hydroxide (0.005-normal) with constant shaking. In order to account for this amphi-electrical behaviour, it is suggested that the potential difference at the surface of colloidal particles is due to adsorption of ions from the solution. Hence the sign depends upon whether cations or anions are in excess in the layers nearest the particles.

The composition and nature of precipitated ferric hydroxide have been made the objects of considerable research. When precipitated from aqueous solutions the hydroxide contains very varying quantities of water, and the problem has been to determine in what manner that water is associated with the complex. This has been attempted in a variety of ways, such as by measuring the rate of dehydration at constant temperature; by studying the dehydration at different temperatures; by determining the vapour pressures during dehydration, etc. A particularly useful method is that of Foote and Saxton, which consists in freezing the precipitate at low temperatures and calculating the quantity of water physically attached to it by observing dilatometrically the change in volume undergone in consequence of the expansion of this water in forming ice. If, in the case of ferric hydroxide, this amount is subtracted from the total amount, as determined by finally heating to redness and weighing the anhydrous ferric oxide, Fe2O3, the difference gives the weight of water chemically combined.

Application of certain of these methods, and particularly the last-named, to precipitated ferric hydroxide, indicates that its contained water is present in three ways, namely: -
  1. Very loosely attached, which readily freezes out at about -5° C.
  2. More intimately associated with the precipitate, and considerably more difficult to remove. It is known as capillary water and requires a temperature of the order of - 30° C. to effect its complete removal by freezing.
  3. Chemically combined. This amount corresponds to the formula Fe2O3.4.25H2O for the hydroxide.

Attempts to determine the molecular weight of colloidal ferric hydroxide lead to very high values. Thus, a colloidal solution prepared by addition of ammonium carbonate to ferric chloride solution was purified by dialysis, and the freezing-point determined of that portion which would not pass through a collodion membrane. The point was only slightly lower than that of the filtrate, indicative of a molecular weight of 3120 for the colloid.

A result of this kind is uncertain, however, owing to the exceedingly small difference in temperature to be registered. Thus Kraft found a colloidal solution containing 3.38 per cent, of ferric hydroxide, and 0.098 per cent, of ferric chloride froze at a temperature within 0.001° C. of the freezing-point of pure water. Other investigators have sometimes obtained negative and at other times positive differences in the freezing- points.

Similar difficulties are encountered when attempts are made to determine the molecular weight by means of measurements of the osmotic pressures of colloidal solutions,5 the observed pressures being exceedingly small, thus indicating a high molecular weight, but one of uncertain value.

For particulars of further researches on colloidal ferric hydroxide the reader is referred to the subjoined references.

The heat of coagulation of colloidal ferric hydroxide with potassium oxalate has been studied by Doerinckel.

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