Chemical Properties of Pure Compact Iron
|When exposed to dry air at ordinary temperatures, iron retains its silver-white appearance. If the air, however, is moist, and the temperature fluctuating so that liquid water collects on the surface of the metal, oxidation or rusting occurs.|
When heated in air or oxygen a piece of polished iron undergoes no apparent change below a temperature of about 150° C. Further heating results in tarnishing. As this is merely slight superficial oxidation, the temperature at which it becomes distinctly visible depends upon the duration of the experiment. Thus, for example, prolonged heating at 170° C. may result in the production of a pale straw colour, although for short periods of time a temperature of 220° C. is normally required to produce the same effect. Given reasonably uniform conditions, however, the extent of the oxidation, which may be judged by the characteristic hues imparted to the iron, is a fairly accurate indication of the temperature. Workmen avail themselves of this with remarkable skill in tempering steels, the data usually accepted being as follows: -
These tempering colours are obtained even in the presence of such dry air as that obtained by continued exposure to phosphorus pentoxide, clearly proving that the reaction is one of direct oxidation of the metal, and therefore entirely distinct from ordinary rusting which involves a preliminary solution of the metal. The oxide produced is usually believed to have the composition represented by the formula Fe3O4. According to Mosander, this is correct in so far as the extreme outer layers of oxide are concerned, those occurring nearer the metal itself having some such formula as Fe2O3.6FeO, or Fe3O4.5FeO. This does not necessarily imply the existence of a definite compound, however; it is more reasonable to assume that ferrous oxide is first formed, and this is relatively slowly converted into ferroso-ferric oxide on account of the difficulty experienced by the oxygen in penetrating the outer layers. Hence the above substance is really a mixture of Fe3O4 and FeO.
When iron wire is strongly heated in an atmosphere of oxygen it burns with a brilliant flame. A pleasing lecture experiment consists in holding a bunch of fine iron wire in the upper part of a Bunsen flame and allowing a jet of oxygen from a gas cylinder to impinge upon the whole.
A modification of this experiment consists in placing a small piece of glowing wood charcoal on a heap of purified iron filings and a stream of oxygen directed upon it. Vigorous combustion ensues, the whole fusing to a white-hot mass of ferroso-ferric oxide, Fe3O4.
According to Charpy, when iron is heated in contact with carbon (graphite) it does not become carburised even at 950° C. unless at least traces of oxygen or an oxide of carbon are present., but this is disputed.
Iron absorbs silicon, when heated with that element, at temperatures considerably below 950° C.
When heated in steam electrolytic iron undergoes no change until about 330° C., when tarnishing begins to take place. At 400° С. a small but measurable quantity of hydrogen is formed, and the velocity of the reaction increases rapidly with further rise of temperature. The reaction appears to take place in three stages, involving
- Dissociation of the steam, H2O ⇔ H2 + O.
- Formation of ferrous oxide, Fe + O ⇔ FeO.
- Oxidation to ferroso-ferric oxide, 3FeO + O ⇔ Fe3O4.
For ordinary iron shavings, the lowest temperature at which hydrogen is evolved is about 300° C., and the optimum yield is obtained at 800° C.
If the reaction is allowed to take place in an enclosed space, it does not proceed to completion. Equilibrium is set up, and the reaction obeys the law of Mass Action. The initial and final stages of the equilibrium may be represented as follows: -
3Fe + 4H2O ⇔ Fe3O4 + 4H2.
Designating the pressure of water vapour as p1 when equilibrium has been reached, and the hydrogen pressure as p2, Preuner obtained the following mean values for the ratio p1/p2 -
When magnetic oxide of iron is heated in a current of hydrogen gas, one of the gaseous phases, namely steam, is swept away, with the result that the oxide is readily reduced to the metal. The same is true for the other oxides of iron, and the reduction has been observed to commence at relatively low temperatures, namely at about 305° C. with magnetic oxide and 370° C. with ferrous oxide. Hilpert finds that if the ferrous oxide has not previously been heated above 400° C., it can be reduced at 280° C. in hydrogen; but if previously heated to 1200° C., reduction is not apparent below 330° C. This shows that the physical condition of the oxide has an important influence upon its dissociation pressure.
The oxidation of iron with steam is used technically as a means of protecting the metal against corrosion. This is the principle of the Bower-Barff process.
With nitrous oxide at 200° C. ferrous oxide is produced. Nitrogen is absorbed by the heated metal to a slight extent, particularly when melted under a high pressure of the gas, yielding the nitride. The nitride is also produced by heating the metal to 800° C. in an atmosphere of ammonia, the physical properties of the metal undergoing considerable alteration.
When iron is heated in contact with carbon and its oxides, many interesting reactions occur. At 900° C. in a current of carbon dioxide iron yields ferrous oxide, whilst at 1200° C. magnetite is produced, which is both magnetic and crystalline. Ignition in carbon monoxide at 1000° C. yields ferrous oxide.
The reaction between carbon monoxide and iron at 650° C. involves the deposition of carbon in those cases where the gas is allowed to pass over the metal in a continuous stream. If, however, the gas and metal are allowed to remain in contact in a closed vessel at 650° C., no carbon is deposited; but a carbide, most probably cementite, Fe3C, is formed, and an oxide. The iron may be dissolved in acid without leaving any carbonaceous residue. What the precise nature of the reactions may be is not certain. The products appear to be the result of many balanced reactions which may be represented by the following equations: -
- 3Fe + 2CO ⇔ Fe3C + CO2.
- 2CO ⇔ C + CO2.
- Fe3C + 4CO2 ⇔ 3FeO + 5CO.
The carbide content falls with increase of pressure of carbon dioxide, becoming nil with a partial pressure of 43 per cent, of carbon dioxide, the remaining gas being carbon monoxide. At 850° C., iron decomposes carbon monoxide, yielding triferro carbide (cementite) and carbon dioxide; thus
3Fe + 2CO = Fe3C + CO2. The iron carbide now decomposes more carbon monoxide, yielding the dioxide and unstable higher carbides of iron, which latter dissociate into free carbon and cementite.
When carbon monoxide and hydrogen saturated with water vapour is passed over iron at 250° to 300° C., several interesting changes may take place. Formaldehyde may be detected in the distillate and a fatty substance, m.pt. 35° to 36° C., isolated by extraction with ether.
These reactions recall the behaviour of nickel under analogous conditions.
Iron is readily corroded by moist chlorine at atmospheric temperatures, and when strongly heated in a current of the dry gas yields ferric chloride, which volatilises and condenses on a cooler part of the apparatus in a beautifully crystalline form.
A small ball of steel wool, if sprinkled with antimony, will ignite in chlorine at the ordinary temperature. Iron is not attacked by anhydrous liquid chlorine.
When heated in hydrogen chloride, iron yields ferrous chloride, free hydrogen being evolved.
Heated to dull redness in bromine vapour, iron yields a yellow crystalline dibromide, FeBr2, or the dark red ferric salt, FeBr3, according to circumstances, excess of bromine vapour being essential to produce the latter compound.
Heated with excess of iodine, iron yields a grey mass of ferrous iodide; the same salt is formed when iron filings are triturated with iodine.
When heated together, iron and sulphur readily unite to yield nonmagnetic ferrous sulphide.
Sulphur dioxide, when dry, has no action on iron, even at 100° C., but the metal is slightly attacked by the moist gas. Liquid sulphur dioxide has likewise but little action under atmospheric pressure, but in refrigerators, where the temperature is liable to rise somewhat during compression of the gas, corrosion of the iron is appreciable.
When heated to 150° to 200° C. in sealed tubes with thionyl chloride, ferric chloride is obtained in accordance with the equation
2Fe + 4SOCl2 = 2FeCl3 + 2SO2 + S2Cl2.
If, however, the metal is present in excess, ferrous sulphide and chloride result. Thus:
3Fe + 2SOCl2 = 2FeCl2 + FeS + SO2.
Nickel is not attacked under these conditions.
When heated with sulphuryl chloride, SO2Cl2, anhydrous ferric chloride is obtained, large crystals being formed in favourable circumstances. Sulphuryl fluoride, SO2F2, is without action on iron, even at red heat.
Silicon tetrachloride vapour is decomposed by iron at high temperatures, yielding diferrosilicide and ferrous chloride,
SiCl4 + 4Fe = Fe2Si + 2FeCl2.
Nitric oxide is reduced by moist iron over mercury, yielding a mixture of nitrous oxide and nitrogen. The nitrous oxide is slowly reduced to nitrogen.
Iron precipitates copper, silver, antimony, lead, and tin from solutions of their salts. In the case of tin the deposit may be exceedingly small, as it forms a thin protecting layer on the surface of the iron. For this reason iron dissolves much less rapidly in an acid if a tin salt is present. A pretty experiment consists in immersing a strip of iron in a tube containing a solution, the bottom half of which consists of a concentrated electrolyte in which some tin salt is dissolved, and the upper half a dilute solution of the electrolyte without any tin. The lower portion of the iron strip becomes covered with crystals of tin, whilst the upper portion is quite free, although if completely immersed in the lower portion no tin deposit would be visible.
In the precipitation of silver from dilute solutions of silver nitrate, unworked iron is found to be more active than the worked or strained metal.
The precipitation of copper from solution by means of iron is used commercially as a wet method of extracting copper from pyrites. The kind of scrap iron used appears to influence very markedly the physical condition in which the copper is precipitated. The mud which collects when neutral copper sulphate solutions are reduced in this manner contains a basic ferric sulphate, Fe2(OH)4.SO4, but addition of a little dilute sulphuric acid prevents this, and enables a clean deposit of copper to be obtained.
Iron dissolves in ammonium persulphate solution, yielding ferrous sulphate, a portion of which undergoes further oxidation.
A solution of bleaching powder rapidly attacks iron, with evolution of oxygen. A concentrated solution of ferric chloride has little action on iron, but the metal readily dissolves in a dilute solution.
Iron has been recommended as an electrode, together with carbon in a galvanic cell, the electrolyte being a concentrated solution of ferric chloride. The iron dissolves, yielding ferrous chloride,
2FeCl3 + Fe = 3FeCl2,
which is converted into the ferric state again by passage of chlorine gas. The generated current, of course, passes in the direction from the carbon to the iron through the connecting wire. The cell may be varied by immersing the carbon in ferric chloride solution in a porous cell, surrounded by a solution of sodium chloride into which the iron electrode is made to dip.
Iron and iodine interact in the presence of water, evolving heat, ferrous iodide passing into solution. The reaction appears to take place in stages involving the formation of ferric iodide, which decomposes into ferric oxide and hydrogen iodide, the last-named attacking the free iron with the formation of ferrous iodide.
The Action of Acids upon Iron
|As a general rule, when acids act upon iron hydrogen gas is evolved and the metal passes into solution in the form of a ferrous salt. Thus, for example, dilute sulphuric acid solution reacts as follows: - |
Fe + H2SO4 = FeSO4 + H2.
Prolonged exposure of such solutions to the air results in the more or less complete conversion of the ferrous salt to the ferric condition, which may either remain in solution or be precipitated as a basic salt according to circumstances. After exposure to acid attack, the undissolved pieces of iron very frequently exhibit a peculiar brittleness, but revert to their original condition if kept for one or two days at ordinary temperatures, or in the course of a few hours if gently warmed. This phenomenon was first observed by Johnson in 1873, and the following year was proved by Reynolds to be entirely due to the occlusion of hydrogen within the pores of the metal.
The purer forms of iron (wrought iron and steel) appear to be much more susceptible to this kind of reaction than cast iron. If the attacking acid is readily reducible by hydrogen, many secondary reactions may take place. Thus with nitric acid, oxides of nitrogen and ammonia may be evolved; whilst with selenic acid a deposit of elementary selenium is obtained (see below). When iron is exposed to the action of acids that are also powerful oxidisers - such as, for example, fairly concentrated solutions of nitric and chromic acids, - it is frequently rendered inert or passive. Its surface may remain perfectly bright, but the metal does not show any appreciable solution in the acid, and if removed and immersed in dilute solutions of such salts as copper and silver sulphates, no reaction takes place, although ordinary active iron would cause an immediate precipitation of the more electronegative metal.
Sulphuric acid, in all stages of dilution, attacks iron. Hydrogen is evolved in the case of the diluted acid, both in the cold and on warming. The concentrated acid yields hydrogen in the cold, but sulphur dioxide is the main gaseous product at 160° C. There is no hydrogen sulphide evolved.
Hydrochloric acid readily dissolves iron, hydrogen being evolved, ferrous chloride passing into solution. The rate of solution appears to be doubled for each rise of 10° C. in temperature for concentrations of acid ranging from 25 to 216 grams per litre. Increase in concentration of the acid likewise accelerates the rate of solution, the velocity being doubled for each increase of 30 grams of hydrogen chloride per litre. The presence of arsenious oxide exerts a remarkably retarding effect upon the activity of the acid.
The action of nitric acid upon iron is both interesting and involved.
When iron is dissolved in acid of density 1.40, both nitrogen peroxide and nitric oxide are obtained in proportions varying with the amount of solvent. It is believed that the nitric oxide is not a primary product of the reaction, but a secondary product formed by reduction of the peroxide.
On reducing the density of the acid to 1.30, nitrous oxide is formed in amount equivalent to 7 per cent, of the iron dissolved. With acid of density 1.25, the reaction is further complicated, nitrogen and ammonia being produced: the former gas reaches a maximum in the case of acid of density 1.15, and then falling, whilst the ammonia reaches a maximum with acid of density 1.05.
In the case of the very dilute acid of density less than 1.034, solution is not accompanied by evolution of gas, but with the formation of ferrous nitrate and ammonium nitrate. If acid of density ranging from 1.034 upwards to 1.115 be employed, the resulting solution consists of a mixture of ferrous nitrate and ferric nitrate, the proportion of the latter increasing with the density of the acid. With acid of greater density than 1.115, ferric nitrate is the main product, and acid of density 1.33 is therefore recommended for the preparation of this salt.
The pure, anhydrous acid, free from nitrous acid and dissolved oxides of nitrogen, is without action on pure qualities of iron even at the boiling-point.
In nitric acid of density 1.45 iron remains bright and refuses to dissolve. When removed and immersed in solutions of copper sulphate, no change takes place. In other words, the metal has been rendered "noble" or "passive."
Carbonic acid acts on iron, yielding ferrous carbonate or soluble ferrous hydrogen carbonate, and evolving hydrogen,
Fe + CO2 + H2O = FeCO3 + H2.
This reaction proceeds slowly but steadily in the cold, and is accelerated by vigorous shaking and the employment of iron filings.
Selenic acid dissolves iron with the production of ferrous selenate, FeSeO4, and a deposit of selenium which collects on the surface of the undissolved metal, thereby greatly retarding the reaction. No hydrogen is evolved, and the selenium is presumably obtained by the action of nascent hydrogen upon the excess of selenic acid. Thus: -
6H + H2SeO4 = 4H2O + Se.
The net result of the reaction may be represented by the equation 3Fe + 4H2SeO4 = 3FeSeO4 + Se + 4H2O.
Chloric acid, both dilute and in concentrated solution, attacks iron, the metal dissolving without evolution of hydrogen, the chloric acid being reduced.
Aqueous hypochlorous acid slowly attacks iron, both hydrogen and chlorine being evolved.
The Action of Alkalies upon Iron
|Dilute solutions of the hydroxides of the alkali and alkaline earth metals are capable of preserving iron for an indefinite period from corrosion, provided they are kept out of contact with carbon dioxide, etc. When exposed to the open air, the alkali is readily neutralised by absorption of carbon dioxide, and the iron begins to corrode. Concentrated solutions of caustic soda or potash exert a slight solvent action upon iron, but they inhibit corrosion even when continuously exposed to the air, because solutions of their carbonates at similar concentrations likewise inhibit corrosion, so that the absorption of carbon dioxide from the air is immaterial. By continued immersion in a saturated solution of sodium hydroxide at 100° C., wrought iron is rendered brittle owing to the absorption of hydrogen. If, however, the immersion is still further prolonged, the metal loses its brittleness. It has been suggested that the initial brittleness is due to the more rapid absorption of hydrogen by the amorphous layers between the crystals of the metal than by the crystals themselves, whereby a certain amount of expansion occurs, forcing the crystals apart and weakening their cohesion. After prolonged immersion, however, the hydrogen has had time to diffuse into the crystals themselves, and thus to reduce the intercrystalline strain so that the brittleness disappears.|
When iron is made the anode in a concentrated solution of caustic soda or potash, a low current density (about 0.001 ampere per sq. centimetre) being employed, the metal is oxidised to the alkali ferrate and passes into solution. Sodium hydroxide acts more rapidly than potassium hydroxide, probably on account of the superior solubility of sodium ferrate.
When electrolytic iron foil is immersed in concentrated solutions of sodium or potassium hydroxide for several weeks, and, after thorough cleaning, allowed to corrode in distilled water, the latter gradually becomes contaminated with traces of sodium or potassium salts, the presence of which can be detected by the spectroscope or by the usual Bunsen flame test. Similar results have been obtained with lithium hydroxide, barium hydroxide and with ammonia. It appears probable that the alkali penetrates in minute quantities into the metal between the ferrite crystals, possibly in consequence of a certain amount of porosity in the intercrystalline cement. This theory is supported by the fact that iron which has been soaked in alkali invariably " pits " badly when allowed to corrode; and pitting is usually associated with electro-chemical activity between the boundaries of the crystals.
On boiling pure iron gently in concentrated sodium hydroxide solution, ferrous oxide passes into solution, and upon exposure to air oxidation to ferric oxide takes place.
Fused caustic soda and potash, particularly under the influence of pressure, attack iron readily.
Between 550° and 660° C. fused caustic potash attacks iron appreciably, but the liberation of hydrogen or potassium or the formation of water has not been demonstrated. What, therefore, the precise action is, remains uncertain.
In the case of fused caustic soda between 400° and 720° C. the evolution of hydrogen and the formation of water have both been observed, indicative of the production of a compound Fe(ONa)x. Since caustic soda can be thoroughly dehydrated at 400° C. and undergoes no further loss on heating to 720° C., the formation of water on heating with iron cannot be due to simple decomposition according to the equation
2NaOH = Na2O + H2O.
At white heat, according to Deville, metallic sodium is obtained.
When fused with sodium peroxide, iron yields dark red tabular crystals of the monohydrate, Fe2O3.H2O. Density 3.8.