Chemical elements
  Iron
    History of Iron
    Mineralogy
    Isotopes
    Energy
    Production
    Application
    Physical Properties
      Allotropy
      Occlusion of Gases
      Absorption of Nascent Hydrogen
      Permeability to Gases
      Passivity of Iron
      Iron Powder
      Iron sponge
      Iron amalgam
      Colloidal Iron
      Pyrophoric Iron
      Catalyst
      Iron Ions
      Atomic Weight
    Chemical Properties
    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

Passivity of Iron






In 1790 Keir drew attention to the fact that a piece of iron, when placed in contact with nitric acid of density 1.45, is rendered inactive or passive. It does not appear to dissolve in the acid; when placed in a dilute solution of copper sulphate it does not effect the deposition of copper; and when immersed in ordinary water it exhibits remarkable resistance to corrosion. And this, in many cases, to quote the words of Keir "without the least diminution of metallic splendour or change of colour".

Nitric acid is not the only medium which may be employed for the passivification of iron. Other acids such as chromic, iodic, arsenic, chloric, hydronitric, etc., exert a similar action, as also do mixtures of two or more acids or acids and certain salts, such, for example, as sulphuric and nitrous acids, or a mixture of these with potassium iodate, etc. Iron may also be rendered passive by immersion in aqueous solutions of many oxidising metallic salts such as silver nitrate, lead nitrate, potassium permanganate, soluble bichromates, etc. Hydrogen peroxide passivifies the metal, and a piece of cleaned electrolytic foil may be immersed in a warm, dilute solution of perhydrol without evincing any sign of corrosion, even although its surface is continuously swept with bubbles of oxygen due to the catalytic decomposition of the peroxide. Similarly iron may be rendered passive by making it the anode in an electrolyte containing water.

Gases, also, may cause iron to assume the passive state. Compressed nitric oxide is a case in point. The vapours of concentrated nitric acid have for many years been known to act similarly.

Dry nitrogen peroxide induces a more intense passivity than nitric acid when allowed to come into contact with iron - an observation which suggests an explanation for the fact that iron is only rendered passive by nitric acid which is either yellow or red, whilst passive iron is actually rendered active by immersion in colourless nitric acid solutions. Apparently only acids of such concentrations as are capable of yielding nitrogen peroxide in contact with the metal are able to exert a passivifying action. This is further supported by the fact that by passing nitrogen peroxide into those concentrations of nitric acid in which iron is normally active, the metal becomes passive.


Testing for Passivity

According to Heathcote, iron may be regarded as passive when no chemical action can be detected by the unaided eye after immersing, shaking, and finally holding motionless a piece of the metal in nitric acid of density 1.20, at the room temperature (15° to 17° C.). This is a preferable method to that of Schonbein, who employed nitric acid of density 1.35 in a similar manner, because this latter concentration of acid is sufficient to render active iron passive, whereas acid of density 1.20 does not do so, at the room temperature.

It is important to note this temperature restriction, however, for whereas nitric acid of density 1.250 does not render iron passive at 0° C., yet if the temperature is raised to 10° C. or above, the metal is readily passivified by it.

Dunstan recommends as a convenient test for passivity the employment of a 0.5 per cent, solution of copper sulphate. This solution at once deposits a film of copper on active iron, whilst the passive metal will remain, ofttimes for hours, bright and apparently entirely unaffected.

The action of distilled water on iron may also be used as a test for passivity. Active iron immersed in it usually shows visible signs of corrosion in from 8 to 10 minutes, whilst the passive metal may remain perfectly bright for an hour or more. When, in this latter case, corrosion begins, the action is very local for a considerable time.

Passivity is also readily detected electrically by the difference in potential between passive and active iron, the former having a lower potential. This is a particularly convenient method of detecting passivity caused by anodic polarisation.

Cause of Passivity

Several theories have been suggested to account for passivity, but no one theory suffices to explain every case. The probability is that several kinds of passivity exist and that not a few of the different theories are correct in certain cases.

The oxide theory of Iron Passivity

The oxide theory, apparently the first to be suggested, postulates the formation of a thin layer of oxide upon the surface of the metal, thereby protecting the underlying portions from reacting. Such an action is not at all uncommon. Thus ordinary lead is well known to resist atmospheric corrosion by reason of the thin protective film of suboxide formed on its surface. Aluminium behaves similarly, and the difficulty of finding a suitable solder for the metal is connected with this oxide layer. The theory receives substantial support from the fact that most passivifying agents are oxidisers, and is in entire harmony with the fact that passivity is a surface phenomenon and can be removed by mechanical processes or by heating in a reducing atmosphere.

The action of 3 per cent, ozone in the cold is a sensitive method of detecting a film of oxide on a metallic surface, inasmuch as a visible deposit of oxide is immediately produced. Iron which has been passivified by immersion in nitric acid or anodically in dilute sulphuric acid instantly reacts with ozone, showing that its surface is coated with oxide. Active iron, on the other hand, exhibits no such sensitiveness towards ozone. The nature of the oxide is uncertain. It has been suggested that it is not likely to be higher than FeO3, but may be FeO2.

Exception to this oxide theory, however, has been taken by Muller and Konigsberger, on the ground that the reflecting power of iron, rendered passive anodically when immersed in alkaline solution, remains undimmed, whereas if a layer of oxide were formed an alteration would be expected. But it is not necessary to postulate the formation of a thick layer of oxide. If of merely molecular dimensions, it would still preserve the underlying metal from attack, whilst a thickness comparable with that of the length of a light wave would be necessary to affect the reflecting power.

On the other hand, when highly polished iron rods are immersed in concentrated nitric acid a visible dulling usually ensues. This is exactly what might be anticipated in accordance with the oxide theory, inasmuch as such a powerful oxidiser might easily induce the formation of a sufficient thickness of oxide to be visible.

Gaseous Film Theory of Iron Passivity

It is well known that gases frequently adhere with extraordinary tenacity in the form of thin films to solid surfaces, and it has been suggested that some such gaseous film is the cause of the passivity assumed by iron in certain circumstances. The film need not necessarily be one of oxygen, although in perhaps the majority of cases this may be the predominant gas. Nitric oxide or nitrogen gas may likewise induce passivity in this manner.

It may be presumed that whilst gaseous films of widely differing gases may in the main behave alike, they will yet show minor differences or eccentricities in their behaviour. An explanation is thus forthcoming for the differences in the behaviour of irons passivified in various ways. Thus, for example, iron that has been rendered passive by anodic polarisation in sulphuric acid differs slightly from that passivified by immersion in concentrated nitric acid. In the former case it is therefore presumed that the metal is covered with a film of gaseous oxygen, but in the latter with a film of some oxide of nitrogen.

This, like the oxide theory, has much to recommend it. An explanation is afforded for the superficial nature of passivity, and the theory also offers an explanation for the passivity induced in certain cases where oxygen cannot be the cause.

A close connection has been found to exist between passivity and photo-electric behaviour. Thus dry, active iron is found to exhibit a considerably higher photo-electric activity than the metal rendered passive by immersion in concentrated nitric acid or by anodic polarisation in dilute sulphuric acid; and this is regarded as supporting the gaseous film theory.

Physical Theories of Iron Passivity

Several theories have been put forward according to which the phenomena of passivity are due to a physical change in the superficial layers of the iron. What the precise nature of that change may be remains a matter of controversy.

According to some authorities, passivity is the normal state of pure iron, the metal only being rendered active by the presence of some catalyst, such as hydrogen or hydrogen ions, which increases its solution pressure. This view receives powerful support from the experiments of Lambert and Thomson, who prepared specimens of exceptionally pure iron by methods already indicated and found them to be remarkably inert or passive.

Other investigators have suggested allotropy as a possible cause, a layer of passive, allotropic metal being formed in contact with the passivifying reagent. This, however, is simply substituting another name for passivity inasmuch as no explanation of this type of allotropy is forthcoming. If this allotropy is due to a rearrangement of the atoms in consequence of change of valency, this theory is equivalent to that of Kriiger and Finkelstein, who regard ordinary active iron as the ferrous or divalent variety, the passive metal being ferric or trivalent. Since valency is an electrical phenomenon, the valency hypothesis is equivalent to assuming an electrical difference between the active and passive metal. This is the basis of Miiller's theory.

An attractive hypothesis is that of Smits, who inclines to the view that iron contains, in addition to uncharged atoms and free electrons, two kinds of ions, α and β, which differ in their reactivity, and which are in equilibrium with one another. Thus: -

α ⇔ β.

The production of the passive state is attributed to a disturbance of this equilibrium. During anodic polarisation of iron the more reactive or α ions dissolve with greater rapidity than equilibrium can be established, with the result that an excess of " noble," inert, or passive β ions collects on the surface of the metal, tending to render it passive. This explains why passivity is purely a surface phenomenon. Hydrogen ions, like halogens, are assumed to catalytically accelerate the conversion of a into β ions until equilibrium is re-established.

When iron is immersed in a passivifying reagent such as nitric acid, not only do the α ions dissolve with great rapidity; but the acid, being a powerful oxidiser, quickly removes any hydrogen ions from the surface of the metal and thus reduces the tendency for equilibrium to be re-established, and the metal again becomes thoroughly passive.

If, now, the current density is increased to such an extent that oxygen begins to be evolved at the anode, all the hydrogen ions, normally present in ordinary iron, will be completely expelled from the surface of the metal, the tendency for reversion to equilibrium of the α and β ions -being reduced to a minimum. In other words, the metal is rendered thoroughly passive.

If, on the other hand, iron is immersed in a solution of ferrous chloride, although the a ions rapidly dissolve passivity is not induced because the chlorine and halogen ions rapidly readjust the equilibrium, and the β ions never accumulate in sufficient excess on the metallic surface to produce passivity.

Iron which has been passivified anodically can be rendered active by the introduction of halogen ions; by the electrolysis of a mixed solution of ferrous sulphate and chloride passivity can be made a periodic phenomenon.

From the foregoing it is evident that widely divergent views are held as to the cause of passivity. For further particulars the reader is referred to the subjoined references.

An instructive series of papers on passivity was read at the general discussion on the subject held by the Faraday Society, November 12, 1913. The papers are printed in the Transactions of the Society for 1914, vol. ix.
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