Iron Salts

Salts of Iron may be roughly divided into two groups, according as the iron behaves as a divalent or trivalent atom. Salts of divalent iron are termed ferrous, and in neutral or faintly alkaline solution are readily oxidised to ferric compounds, in which the iron has a valency of three. Ferrous salts, when hydrated, are usually greenish in colour, copperas or ferrous sulphate being typical. Anhydrous ferrous salts are white or pale yellow in colour, as witness ferrous chloride and bromide. Ferric salts, when crystalline, are also white and opaque (ferric sulphate), colourless (ferric nitrate), or yellow (ferric chloride), although many basic compounds are brown. When in combination with other salts, various colours may appear - as, for example, in the case of iron alum, the delicate violet hue of which is well known. Iron salts impart a distinct and bitter taste to water, one part of iron per million of water being distinctly perceptible to the average individual.

A peculiar property of ferrous salts consists in their power, when in solution, of uniting with nitric oxide, the limit of combination being reached with one molecule of NO to one atom of iron. The substances thereby produced are very unstable, partaking of the nature of additive or associated compounds. Two of these compounds have been isolated, namely, FeSO₄.NO and FeHPO₄.NO; and several other salts, such as FeCl₂.NO and FeBr₂.NO, have been shown to exist in solution. In this respect ferrous salts closely resemble salts of divalent copper, which yield, with nitric oxide, additive compounds of the type CuX₂.NO. Ferric salts - for example, ferric chloride - also combine with nitric oxide to yield unstable compounds. Both ferrous and ferric salts readily decompose barium peroxide, the former after being first oxidised to ferric. The reaction, in the case of ferric chloride, takes place according to the equation: -

6BaO₂ + 4FeCl₃ + 6H₂O = 3O₂ + 4Fe(OH)₃ + 6BaCl₂.

Schonbein, in 1857, drew attention to the fact that ferrous salts are capable of acting as oxygen-carriers in certain circumstances,7 but it was not until thirty-seven years had elapsed that attention was again drawn to the subject by the observation of Fenton that hydrogen peroxide oxidises tartaric acid to dihydroxymaleic acid in the presence of a ferrous salt - in particular, ferrous sulphate. He explained this on the assumption that the divalent iron replaces the two non-hydroxylic hydrogen atoms of tartaric acid, and, upon oxidation by the peroxide to the trivalent condition, breaks away from them. Thus



Suggestive as this theory is, there are other reactions in which ferrous sulphate acts as a catalyser, in which the above explanation cannot possibly hold, and several other theories have been offered. Manchot and Wilhelms believe that a highly oxidised compound of iron, such as Fe₂O₅ or Fe₂O₃.2H₂O₂, is formed by contact of the ferrous salt and peroxide, and that this is reduced by the oxidisable substance. Brode reached a somewhat similar conclusion, suggesting that the peroxide and ferrous salt unite to form an intermediate, highly oxidised compound, whilst Mummery combines the essential features of these two theories into a very attractive hypothesis, according to which ferrous sulphate and hydrogen peroxide unite to form ferrous sulphate perhydrol,



He arrives at this in the following manner. In view of the retention of a molecule of water by the isomorphous members of the ferrous sulphate series, they may be regarded as hemisulphate hemihydrols of the type



With hydrogen peroxide, therefore, ferrous sulphate reacts as follows: -



This perhydrol bears a resemblance in constitution to Caro's per mono - sulphuric acid, H₂SO₅, or



and may be expected to possess the properties of a powerful oxidiser. The reason why such a compound acts more powerfully than hydrogen peroxide itself is attributed to the fact that it is an electrolyte, whereas hydrogen peroxide, for all practical purposes, is not. This perhydrol is alternately produced from, and reconverted into, ferrous sulphate, when an oxidisable substance is present, together with hydrogen peroxide.

Methyl, ethyl, and propyl alcohols are oxidised by permanganate or hydrogen peroxide in the presence of ferrous salts. If ferrous sulphate is employed, the ethyl alcohol is, in dilute solution, oxidised by the permanganate to aldehyde; but in the presence of ferrous oxalate the oxidation proceeds further, acetic acid resulting. Thus

CН₃.СН₂ОН → СН₃.СНО → СН₃.СООН.

These reactions proceed so regularly that they may be followed up quantitatively. Ferric salts have no catalytic influence upon these reactions. Iron salts, however, can act as oxygen-carriers in the absence of such powerful oxidisers as hydrogen peroxide and potassium permanganate. Thus, for example, it is well known that, upon exposure to sunlight, iodine is ordinarily liberated from a solution of mercuric iodide in potassium iodide. Curiously enough, if traces of iron salts are rigidly excluded, the liberation of iodine does not take place.

Ferrous salts accelerate the oxidation of sulphurous acid to sulphuric acid in the presence of oxygen.

Ferrous chloride accelerates the oxidation of stannous chloride solution in air, the maximum effect being obtained with one molecule of FeCl₂ to 100 molecules of SnCl₂.

Ferrous salts accelerate markedly the reaction between persulphates and iodides, as represented, in the case of the potassium salts, by the equation

K₂S₂O₈ + 2KI = 2K₂SO₄ + I₂

In acid solution chlorates are reduced to chlorides by soluble iodides. For example, in the case of the potassium salts the reaction proceeds as follows: -

KClO₃ + 6KI + 3H₂SO₄ = KCl + 3K₂SO₄ + 3I₂ + 3H₂O.

Addition of a small quantity of a ferric or ferrous salt greatly accelerates the reaction, due to the alternate formation of ferrioiodide and reduction to the ferrous salt as follows: -

KClO₃ + 6FeI₂ + 6KI + 3H₂SO₄ = KCl + 3K₂SO₄ + 6FeI₃ + 3H₂O,
6FeI₃ = 6FeI₂ + 3I₂.

It is well known that a mixed solution of ordinary mercuric chloride and ammonium oxalate undergoes decomposition when exposed to light, mercurous chloride being precipitated and carbon dioxide evolved. The reaction may be represented by the equation

2HgCl₂ + (NH₄)₂C₂O₄ = 2HgCl + 2NH₄Cl + 2CO₂.

This change, which proceeds at a measurable rate, has been utilised in quantitative determinations of the intensity of light - so-called actinometric measurements. The sensitiveness of the reaction, however, is largely dependent upon the purity of the salts, traces of iron salts increasing the sensitiveness in proportion to the amount of iron - provided the amount is very small. Indeed, there is reason to believe that if the solution were entirely free from iron no photochemical effect would be observed.

Hydrochloric acid is without action on gold, but addition of a small quantity of ferric chloride causes the gold to dissolve, the ferric salt presumably acting as a chlorine-carrier in the presence of hydrochloric acid and oxygen. Salts of iron are frequently used in organic chemical processes as halogen-carriers.

In certain cases small quantities of ferrous salts act as retarding agents in chemical reactions, and may therefore be termed negative catalysts. For example, ferrous sulphate has long been known to hinder the action of nitric acid on metals.

Traces of ferric salts retard the dissolution of mercury in nitric acid, a phenomenon which is attributed to decomposition of the nitrous acid by the catalytic alternate reduction and oxidation of the iron radicle.

Ferrous salts are readily converted into ferric derivatives in a variety of ways. Thus their solutions are gradually oxidised upon exposure to air, with the deposition of basic ferric salts. The rate of oxidation by air in the presence of free acids is, in the case of ferrous sulphate, proportional to the partial pressure of the oxygen. Hence the addition of inert soluble salts, such as the chlorides and sulphates of sodium, potassium, or magnesium, to the solution reduces the rate of oxidation in proportion as they decrease the solubility of the oxygen.

At 60° C. the relative rates of oxidation of ferrous chloride, sulphate, and acetate, are as follow: - 1: 10: 100. The oxidation appears to depend upon the un-ionised portion of the dissolved salt.

In alkaline solution oxidation of ferrous iron is fairly rapid, but certain acids retard the reaction. Ferrous sulphate, for example, in the presence of free sulphuric acid, is very stable in air. Concentrated hydrochloric acid assists the oxidation, as also do traces of certain substances, such as platinie and cupric chlorides, palladium nitrate, etc.

The aerial oxidation of solutions of ferrous salts may be accelerated by certain micro-organisms, known as iron bacteria. Mumford10 describes an organism through the agency of which a dilute solution of ferrous ammonium sulphate was completely oxidised to ferric hydroxide in fewer than thirty-six hours at 37° C., no iron remaining in solution. There can be no doubt that the natural deposits of bog iron ore, occurring in Sweden and elsewhere, owe their existence to the action of these lowly organisms.

The presence of platinum also appears to accelerate the oxidation of solutions of ferrous salts.

The action of hydrogen peroxide upon ferrous salts is interesting in view of the possible connection between ferrous salts acting as oxygen- carriers in the blood. The decomposition of the peroxide is hindered or " poisoned " by the presence of arsenious oxide, hydrogen sulphide, carbon monoxide, and other well-known poisons, and the subject is worthy of careful consideration.

The usual oxidising media, such as permanganates, bichromates, etc., react instantaneously with ferrous salts, yielding in acid solution the normal ferric salts. Many methods for the quantitative determination of iron are based on these reactions. Thus, for example, with potassium permanganate the oxidation of ferrous sulphate proceeds as follows: -

10FeSO₄ + 2KMnO₄ + 8H₂SO₄ = 5Fe₂(SO₄)₃ + K₂SO₄ + 2MnSO₄ + 8H₂O.

The persistence of the pink colour of the permanganate indicates with great accuracy when sufficient of this reagent has been added. With potassium bichromate the reaction proceeds according to the equation

6FeSO₄ + K₂Cr₂O₇ + 7H₂SO₄ = K₂SO₄ + 3Fe₂(SO₄)₃ + Cr₂(SO₄)₃ + 7H₂O,

the end point of the reaction being determined by removing a drop of the mixture and testing with a solution of potassium ferricyanide, which remains colourless if no ferrous salt is present, but otherwise yields a deep blue colour.

The mechanism of the oxidation of ferrous salts in the various ways mentioned above has been studied by Manchot, who concludes that the first action is to produce a high and generally unstable oxide, termed a primary oxide, which then decomposes into the final oxidation product (that is, the ferric compound) and active oxygen, which latter then oxidises a further portion of the ferrous salt to ferric.

Owing to their ready oxidisability, ferrous, salts are frequently employed as mild reducing agents. Thus ferricyanides are reduced to ferrocyanides by ferrous sulphate in alkaline solution.

In photography ferrous sulphate is used for the reduction of auric chloride, metallic gold being precipitated. Thus

AuCl₃ + 3FeSO₄ = Au + FeCl₃ + Fe₂(SO₄)₃.

On adding a solution of a ferrous salt to an ammoniacal solution of a cupric salt, ferric hydroxide is precipitated, the cupric salt being reduced to cuprous and remaining in solution. This affords a convenient method of preparing ammoniaeal cuprous solutions for the absorption of carbon monoxide in gas analysis.

Ferrous salts reduce nitrites to nitric oxide. On addition, for example, of barium nitrite to ferrous sulphate, barium sulphate is precipitated, and the liquid turns brown. Ferric hydroxide and a basic ferric nitrate are next precipitated, nitric oxide being evolved. Apparently the first product of the reaction is ferrous nitrite, which then spontaneously decomposes in accordance with the equation

6Fe(NO₂)₂ = 10NO + Fe₂O₃ + 2Fe₂O₃.N₂O₅.

Mercuric chloride solution is reduced by ferrous chloride under the influence of light, mercurous chloride being precipitated.

HgCl₂ + FeCl₂ = FeCl₃ + HgCl.

If the relative proportion of mercuric chloride is small, the rate of reduction is almost independent of the concentration of the ferrous salt, and the sensitiveness to light increases with dilution. Equimolecular solutions, however, do not vary to the same extent, but their maximum sensitiveness occurs at a concentration of 3 gram-molecules of each salt per litre. The presence of oxygen is without appreciable influence on the equimolecular solutions.

The reverse reactions, namely conversion of ferric salts into ferrous, are likewise easily effected by means of the usual reducing agents, such as nascent hydrogen, sulphur dioxide, etc. By the introduction of zinc into an acidified solution of a ferric salt, reduction is rapidly caused. Excess of acid slightly retards the reaction. This affords a convenient method of volumetrically determining the presence of ferric iron, the solution after reduction being titrated with permanganate.

Zinc dust is particularly rapid in its action, even in neutral solution. In this latter case, however, the iron is partially precipitated as ferric hydroxide. The reduction takes place with ease in absolute alcohol, and is entirely independent of the presence of occluded hydrogen in the zinc. It would appear, therefore, that the zinc acts directly as a dechlorinator, and that the reduction is not effected by nascent hydrogen. Thus: -

2FeCl₃ + Zn = 2FeCl₂ + ZnCl₂.

Sulphur dioxide may be passed into a solution of a ferric salt for a similar purpose, or it may be generated in the solution by addition of an alkali sulphite and a little dilute mineral acid. Thus, ferric sulphate is reduced in accordance with the equation

Fe₂(SO₄)₃ + SO₂ + 2H₂O = 2FeSO₄ + 2H₂SO₄.

The reducing action of thiosulphates on ferric salts is interesting, ranking as one of the few tetramolecular reactions that have been studied. The reduction proceeds in accordance with the ionic equation

2Fe^••• + 2S₂O₃'' = 2Fe^•• + S₄O₆''.

Reduction of ferric salts with potassium iodide is usually regarded as taking place as follows, in the case of ferric chloride: -

2FeCl₃ + 2KI ⇔ 2FeCl₂ + 2KCl + I₂;

or with ferric sulphate: -

Fe₂(SO₄)₃ + 2KI ⇔ 2FeSO₄ + K₂SO₄ + I₂.

It is quite possible, however, that intermediate iodo-compounds are first formed, and, being unstable, rapidly decompose, liberating iodine. Thus:

2FeCl₂.Cl + 2KI = 2FeCl₂I + 2KCl, 2FeCl₂I = 2FeCl₂ + I₂,

and

(FeSO₄)₂SO₄ + 2KI = 2FeSO₄I + K₂SO₄,
2FeSO₄I = 2FeSO₄ + I₂.

A useful reducing agent is stannous chloride, which reacts with ferric chloride as follows: -

2FeCl₃ + SnCl₂ = 2FeCl₂ + SnCl₄,

the reaction being one of the third order or trimolecular and not bi-molecular, as Kahlenberg suggested, in neutral solution. In the presence of acid the reaction apparently approximates to one of the second order, but the role played by the acid is not clear.

Ferric salts are reduced by organic substances under the influence of light. Thus an alcoholic solution of ferric chloride, when exposed to sunlight, is converted into ferrous chloride, the alcohol being oxidised to aldehyde. For example, in the case of methyl and ethyl alcohols: -

2FeCl₃ + CH₃OH = 2FeCl₂ + H.CHO + 2HCl,
2FeCl₃ + CH₃.CH₂OH = 2FeCl₂ + CH₃.CHO + 2HCl.

In this reaction the light acts, not as a catalyst, but as a source of energy. The quotient



is nearly constant, from which it is concluded that the amount of light energy required to reduce a definite quantity of ferric chloride is nearly constant and independent of the concentration.

Similarly a ferric salt of an organic acid is reduced by sunlight, ferric oxalate being a well-known example.

Fe₂(C₂O₄)₃ = 2FeC₂O₄ + 2CO₂.

In this case the progress of the reaction may be followed by noting the rate of evolution of carbon dioxide. The blue prints used by engineers are prepared by exposing to light, under a drawing which serves as negative, sheets of paper previously soaked in a solution of an organic ferric salt. After a suitable exposure the paper is washed with a solution of potassium ferricyanide, which gives the characteristic Turnbull's blue colour where the light has effected the reduction of the ferric salt.

Ultraviolet light, such as that, for example, emitted by an electric spark passing between aluminium terminals, is very active in reducing ferric salts, its activity being greatly increased by the presence ox sugar. Alteration of temperature appears to exert only a small influence.

It is interesting to note that ferric salts of organic acids such as citric, tartaric, etc., are not as a rule ferric salts in the ordinary acceptation of the term. The iron has entered into the electro-negative radicle in an analogous manner to copper in the organic copper derivatives. In ferric oxalate and in the ferricyanides, which latter do not contain hydroxylic or carboxylic groups, the iron is similarly in the negative radicle.

Ferrous oxalate is only slightly soluble in water, but the double salt K₂Fe(C₂O₄)₂.2H₂O dissolves more readily. The iron is mainly present in solution as the complex anion Fe(C₂O₄)₂'', but this is not very stable, owing to dissociation, which takes place as follows: -

Fe(C₂O₄)₂ ⇔ FeC₂O₄(solid) + C₂O₄'',

unless an excess of alkali oxalate is present, ferrous oxalate being precipitated.

The magnetic properties of iron salts have been the subject of a considerable amount of investigation.

Both ferrous and ferric salts can be reduced to the metal in a variety of ways. In solution many of them are reduced by the introduction of more electro-positive metals such as magnesium, zinc, or aluminium, as also by electrolysis. In the dry way they are reduced by heating in a current of hydrogen or carbon monoxide, or with zinc dust or by ignition with carbon in an air blast.

When ferric acetate solution is exposed to hydrogen gas at 350° C. and under 230 atmospheres pressure, ferric oxide is precipitated as an anhydrous red mass, insoluble in water but soluble in hydrochloric acid. At 400° C. and under a pressure of 420 atmospheres, metallic iron is obtained.