Chemical elements
  Iron
    History of Iron
    Mineralogy
    Isotopes
    Energy
    Production
    Application
    Physical Properties
    Chemical Properties
      Iron Hydride
      Ferrous fluoride
      Aluminium pentafluoferrite
      Ferric fluoride
      Ammonium ferrifluoride
      Barium ferrifluoride
      Potassium ferrifluoride
      Sodium ferrifluoride
      Thallous ferrifluoride
      Ferrous diferrifluoride
      Ferrous monoferrifluoride
      Ferrous chloride
      Ammonium tetrachlorferrite
      Ferric chloride
      Tetrachlorferrates
      Pentachlorferrates
      Ferroso-ferric chloride
      Ferrous perchlorate
      Ferric perchlorate
      Ferrous chlorate
      Ferric chlorate
      Ferrous Oxychlorides
      Ferrous bromide
      Ferric bromide
      Ferric chloro-bromide
      Ferrous bromate
      Ferrous iodide
      Ferric iodide
      Ferric iodate
      Ferrous oxide
      Ferrous hydroxide
      Triferric tetroxide
      Ferric oxide
      Ferrous acid
      Calcium ferrite
      Cobalt ferrite
      Cupric ferrite
      Cuprous ferrite
      Magnesium ferrite
      Nickel ferrite
      Potassium ferrite
      Sodium ferrite
      Zinc ferrite
      Barium ferrate
      Strontium ferrate
      Barium perferrate
      Calcium perferrate
      Potassium perferrate
      Sodium perferrate
      Strontium perferrate
      Iron Subsulphides
      Ferrous sulphide
      Ferric sulphide
      Potassium ferric sulphide
      Sodium ferric sulphide
      Cuprous ferric sulphide
      Iron disulphide
      Ferrous sulphite
      Ferric sulphite
      Potassium ferri-tetrasulphite
      Potassium ferri-disulphite
      Potassium ferri-sulphite
      Ammonium ferri-sulphite
      Sodium ferri-disulphite
      Sodium hydrogen ferri-tetrasulphite
      Ferrous sulphate
      Ferrous copper sulphate Fe
      Ferrous ammonium sulphate
      Ferrous potassium sulphate
      Ferrous aluminium sulphate
      Basic ferrous sulphate
      Ferric sulphate
      Ammonium ferri-disulphate
      Trisodium ferri-trisulphate
      Ferric Alums
      Ferric ammonium alum
      Ferric potassium alum
      Ferric rubidium alum
      Ferroso-ferric sulphate
      Ferrous amido-sulphonate
      Ferric amido-sulphonate
      Ferrous thiosulphate
      Ferrous pyrosulphate
      Ferrous tetrathionate
      Ferric selenide
      Iron diselenide
      Iron Selenites
      Ferrous selenate
      Ferric rubidium selenium alum
      Ferric caesium selenium alum
      Ferric tellurite
      Ferrous chromite
      Ferrous chromate
      Iron nitride
      Nitro-Iron
      Ferrous nitrate
      Ferric nitrate
      Ferrous Nitroso Salts
      Potassium ferro-heptanitroso sulphide
      Sodium ferro-heptanitroso sulphide
      Ammonium ferro-heptanitroso sulphide
      Tetramethyl ammonium ferro-heptanitroso sulphide
      Ferro-dinitroso Sulphides
      Potassium ferro-dinitroso thiosulphate
      Triferro phosphide
      Diferro phosphide
      Iron monophosphide
      Iron sesqui-phosphide
      Ferrous hypophosphite
      Ferric hypophosphite
      Ferrous phosphite
      Ferric phosphite
      Ferrous orthophosphate
      Ferrous hydrogen orthophosphate
      Ferrous dihydrogen orthophosphate
      Ferric orthophosphate
      Sodium ferri-diorthophosphate
      Ammonium ferri-diorthophosphate
      Sodium ferri-triorthophosphate
      Ferric dihydrogen orthophosphate
      Acid ferric orthophosphate
      Ferrous metaphosphate
      Ferric metaphosphate
      Ferrous pyrophosphate
      Ferric pyrophosphate
      Hydrogen ferri-pyrophosphate
      Sodium ferro-pyrophosphate
      Ferrous thio-orthophosphite
      Ferrous thio-orthophosphate
      Ferrous thio-pyrophosphite
      Ferrous thio-pyrophosphate
      Iron sub-arsenide
      Iron mon-arsenide
      Iron sesqui-arsenide
      Iron di-arsenide
      Iron thio-arsenide
      Ferrous met-arsenite
      Ferric arsenite
      Ferrous ortho-arsenate
      Ferric ortho-arsenate
      Ferro mono-antimonide
      The di-antimonide
      Ferrous thio-antimonite
      Ferric ortho-antimonate
      Triferro carbide
      Diferro carbide
      Iron dicarbide
      Iron pentacarbonyl
      Diferro nonacarbonyl
      Iron tetracarbonyl
      Ferrous carbonate
      Ferrous bicarbonate
      Ferrous potassium carbonate
      Complex Iron Carbonates
      Ferrous thiocarbonate
      Ferrous thiocarbonate hexammoniate
      Ferrous cyanide
      Ferro-cyanic acid
      Aluminium ferrocyanide
      Aluminium ammonium ferrocyanide
      Ammonium ferrocyanide
      Barium ferrocyanide
      Calcium ferrocyanide
      Calcium ammonium ferrocyanide
      Cobalt ferrocyanide
      Copper ferrocyanide
      Ammonium cuproferrocyanide
      Barium cuproferrocyanide
      Lithium cuproferrocyanide
      Magnesium cuproferrocyanide
      Potassium cuproferrocyanide
      Sodium cuproferrocyanide
      Ammonium cupriferrocyanide
      Potassium cupriferrocyanide
      Potassium ferrous cupriferrocyanide
      Sodium cupriferrocyanide
      Strontium cupriferrocyanide
      Lithium ferrocyanide
      Magnesium ferrocyanide
      Magnesium ammonium ferrocyanide
      Manganese ferrocyanide
      Nickel ferrocyanide
      Potassium ferrocyanide
      Potassium aluminium ferrocyanide
      Potassium barium ferrocyanide
      Potassium calcium ferrocyanide
      Potassium cerium ferrocyanide
      Potassium magnesium ferrocyanide
      Potassium mercuric ferrocyanide
      Silver ferrocyanide
      Sodium ferrocyanide
      Sodium cerium ferrocyanide
      Strontium ferrocyanide
      Thallium ferrocyanide
      Zinc potassium ferrocyanide
      Ferricyanic acid
      Ammonium ferricyanide
      Barium ferricyanide
      Barium potassium ferricyanide
      Calcium ferricyanide
      Calcium potassium ferricyanide
      Cobalt ferricyanide
      Copper ferricyanide
      Lead ferricyanide
      Magnesium ferricyanide
      Mercuric ferricyanide
      Mercurous ferricyanide
      Potassium ferricyanide
      Sodium ferricyanide
      Strontium ferricyanide
      Zinc ferricyanide
      Ferrous hydrogen ferrocyanide
      Ferrous potassium ferrocyanide
      Prussian Blues
      Ferrous ferrocyanide
      Ferric ammonium ferrocyanide
      Nitroprussic acid
      Sodium nitroprusside
      Ammonium nitroprusside
      Barium nitroprusside
      Cobalt nitroprusside
      Nickel nitroprusside
      Potassium nitroprusside
      Carbonyl Penta-Ferrocyanides
      Carbonyl ferrocyanic acid
      Barium carbonyl ferrocyanide
      Copper carbonyl ferrocyanide
      Ferric carbonyl ferrocyanide
      Potassium carbonyl ferrocyanide
      Silver carbonyl ferrocyanide
      Sodium carbonyl ferrocyanide
      Strontium carbonyl ferrocyanide
      Uranyl carbonyl ferrocyanide
      Sodium ammonio ferrocyanide
      Potassium aquo ferrocyanide
      Potassium aquo ferricyanide
      Sodium aquo penta-ferricyanide
      Potassium sulphito ferrocyanide
      Ferrous thiocyanate
      Ferric thiocyanate
      Sodium ferrothiocyanate
      Sodium ferrithiocyanate
      Potassium ferrithiocyanate
      Iron subsilicide
      Iron monosilicide
      Iron disilicide
      Triferro disilicide
      Ferrous orthosilicate
      Ferrous magnesium orthosilicate
      Ferrous metasilicate
      Ferric silicate
      Diferro boride
      Iron monoboride
      Iron diboride
      Ferrous chlorborate
      Ferrous bromborate
    Corrosion
    Iron Salts
    PDB 101m-1aeb
    PDB 1aed-1awd
    PDB 1awp-1beq
    PDB 1bes-1c53
    PDB 1c6o-1ci6
    PDB 1cie-1cry
    PDB 1csu-1dfx
    PDB 1dgb-1dry
    PDB 1ds1-1e08
    PDB 1e0z-1ehj
    PDB 1ehk-1f5o
    PDB 1f5p-1fnp
    PDB 1fnq-1fzi
    PDB 1g08-1gnl
    PDB 1gnt-1h43
    PDB 1h44-1hdb
    PDB 1hds-1i5u
    PDB 1i6d-1iwh
    PDB 1iwi-1jgx
    PDB 1jgy-1k2o
    PDB 1k2r-1kw6
    PDB 1kw8-1lj0
    PDB 1lj1-1m2m
    PDB 1m34-1mko
    PDB 1mkq-1mun
    PDB 1muy-1n9x
    PDB 1naz-1nx4
    PDB 1nx7-1ofe
    PDB 1off-1p3t
    PDB 1p3u-1pmb
    PDB 1po3-1qmq
    PDB 1qn0-1ra0
    PDB 1ra5-1rxg
    PDB 1ry5-1smi
    PDB 1smj-1t71
    PDB 1t85-1u8v
    PDB 1u9m-1uyu
    PDB 1uzr-1vxf
    PDB 1vxg-1wri
    PDB 1wtf-1xlq
    PDB 1xm8-1y4r
    PDB 1y4t-1ygd
    PDB 1yge-1z01
    PDB 1z02-2a9e
    PDB 2aa1-2azq
    PDB 2b0z-2boz
    PDB 2bpb-2ca3
    PDB 2ca4-2cz7
    PDB 2czs-2dyr
    PDB 2dys-2ewk
    PDB 2ewu-2fwl
    PDB 2fwt-2gl3
    PDB 2gln-2hhb
    PDB 2hhd-2ibn
    PDB 2ibz-2jb8
    PDB 2jbl-2mgh
    PDB 2mgi-2o01
    PDB 2o08-2ozy
    PDB 2p0b-2q0i
    PDB 2q0j-2r1h
    PDB 2r1k-2spm
    PDB 2spn-2vbd
    PDB 2vbp-2vzb
    PDB 2vzm-2wiv
    PDB 2wiy-2xj5
    PDB 2xj6-2ylj
    PDB 2yrs-2zon
    PDB 2zoo-3a17
    PDB 3a18-3aes
    PDB 3aet-3bnd
    PDB 3bne-3cir
    PDB 3ciu-3dax
    PDB 3dbg-3e1p
    PDB 3e1q-3eh4
    PDB 3eh5-3fll
    PDB 3fm1-3gas
    PDB 3gb4-3h57
    PDB 3h58-3hrw
    PDB 3hsn-3ir6
    PDB 3ir7-3k9y
    PDB 3k9z-3l4p
    PDB 3l61-3lxi
    PDB 3lyq-3mm8
    PDB 3mm9-3n62
    PDB 3n63-3nlo
    PDB 3nlp-3o0f
    PDB 3o0r-3p6o
    PDB 3p6p-3prq
    PDB 3prr-3sel
    PDB 3sik-3una
    PDB 3unc-4blc
    PDB 4cat-4erg
    PDB 4erm-4nse
    PDB 4pah-8cat
    PDB 8cpp-9nse

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: -

Colour.Temperature. °C.
Pale yellow220
Straw230
Golden yellow243
Brown255
Brown purple265
Purple277
Bright blue288
Full blue293
Dark blue316


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
  1. Dissociation of the steam, H2OH2 + O.
  2. Formation of ferrous oxide, Fe + OFeO.
  3. Oxidation to ferroso-ferric oxide, 3FeO + OFe3O4.

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 + 4H2OFe3O4 + 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 -

Temperature. °C.p1/p2
9000.69
10250.78
11500.86


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: -

  1. 3Fe + 2COFe3C + CO2.
  2. 2COC + CO2.
  3. 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.
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