10.8. Chemistry of Iron Fe, Z=26, 1s²2s²2p⁶3s²3p⁶3d⁶4s² 

data comparison of iron with the other members of the 3d–block and transition metals

Extended data table for IRON

• Uses of IRON

  • Iron is an extremely useful silvery–white magnetic metal but moderately reactive towards moist air (corrodes to rust) and readily dissolves in acids.

  • Apart from cast iron, most iron is converted into steel alloys for many purposes.

  • A steel alloy consists mainly of iron mixed with controlled amounts of carbon C, and other metals like chromium Cr, tungsten W, nickel Ni etc.

  • Carbon steels like mild steel (0.1% – 04% carbon) are used for enumerable objects like car bodies, tin cans, nuts/bolts and piping.

  • Stainless steel, an alloy with chromium, has extremely good anti–corrosion properties.

  • Stainless steel (e.g. Fe + Cr + C + Ni) is used for cutlery and surgical instruments

  • Tungsten (Fe + C +W) and manganese steels (Fe + C + Mn) are very tough and hard wearing and used for cutting tools and high speed drill bits.

  • Invar alloy (Fe + C + Ni) is used in watch making.

  • Wrought iron is tough malleable and ductile and good material for a blacksmith to work with.

  • Cast iron, despite being brittle, is used for manhole covers, guttering, machinery frames and drainpipes.

  • Iron is used in the making electromagnets.

  • Iron(II) sulphate, FeSO₄, is used in medicines and vitamin tablets as a source of iron.

  • Iron(III) oxide, Fe₂O₃, is used in paints, dyes and pigments (the 'red ochre' of prehistoric man!)

• Biological role of iron

  • Iron is an essential element in our diet and is needed for the production of haemoglobin.

  • The iron atom in the haemoglobin molecule helps co–ordinate of the oxygen molecule and hence the transportation of oxygen around the body to the cells of all the tissues.

  • Iron deficiency causes anaemia.

  • Plants require iron for the synthesis of chlorophyll.

• An iron/iron(III) oxide mixture is used as a catalyst in the Haber synthesis of ammonia from hydrogen and nitrogen.

  • N_2(g) + 3H_2(g) ==Fe/Fe₂O₃==> 2NH_3(g)

  • Other catalysis examples:

    • The iron(II)/iron(III) ion catalysis of the oxidation of iodide ions by peroxodisulfate ions is described under homogeneous catalysis in Appendix 6.

• Apart from being a member of the 3d block, iron is a true member of the first transition metal series.

  • The two most common ions, iron(II) Fe²⁺ and iron(III) Fe³⁺ both have electron configurations that include a partly filled 3d sub–shell.

    • This is the most important criteria for considering whether such an element is a transition metal

    • The electron configurations are: Fe²⁺ is [Ar]3d⁶ and Fe³⁺ is [Ar]3d⁵

    • For both iron ions, there is at least one electron that can be promoted to a higher level when the 3d sub–shell is split as the central metal ion interacts and bonds with ligands.

      • Visible light photons absorbed, colour results!

    • For more details see Appendix 4. Electron configuration & complex ion colour theory

The extraction of iron and steelmaking

• BUT for A level, note the oxidation state changes in the process:

  • haematite iron ore is Fe₂O₃ (+3) ==> Fe (0);  carbon C (0) => CO₂ (+4) ==> CO (+2) ==>  CO₂ (+4)

  • diiron(III)iron(II) oxide, magnetite Fe₃O₄ is another ore used as a source of iron.

    • note that Fe₃O₄ = 2Fe³⁺ Fe²⁺ 4O^2–, so iron(II) +2 and iron(III) +3 ions ==> Fe(0)

• Starting with impure iron from blast furnace, the molten iron contains many other elements and the iron is too brittle initially, so there is a need to reduce C and remove others like S and P.

  • This is achieved by the Basic Oxygen Steel making process (BOS) which involves many redox reactions. It is a 'batch process' and can't be used as a continuous production line like iron from the blast furnace.

  • Sulphur is removed early in the process using magnesium:

    • Mg + S ==> MgS (the magnesium sulphide becomes part of slag mixture).

  • C, P, Si and others oxidised by molecular oxygen before scrap iron/steel introduced.

    • e.g. C + O₂ ==> CO₂,   4P + 5O₂ ==> P₄O₁₀,  Si + O₂ ==> SiO₂ 

  • After the oxygen blow the basic oxides CaO/MgO are added to form slag salts with the weak acidic oxides of Si and P, carbon dioxide gas will 'escape' from the mixture, since any calcium carbonate formed would thermally decompose at the high temperature of the furnace.

    • e.g. CaO + SiO₂ ==> CaSiO₃ (calcium silicate)

      • 6CaO + P₄O₁₀ ==> 2Ca₃(PO₄)₂ (calcium phosphate(V), forms part of slag)

    • The oxides of Mn/Fe also collect in the slag, so some iron is wasted and the Mn might be added in a controlled way later for a particular steel specification.

    • The toxic carbon monoxide formed must be dealt with and not allowed out into the atmosphere, it can be burned as a fuel to harmless carbon dioxide.

  • It is important to keep track of temperature and composition by thermocouple probe and atomic emission spectroscopy.

  • The elements are oxidised in a sequence in exothermic reactions (no extra heat needed), so temperature control is essential to avoid wasting energy and converter lining damage.

  • The added scrap iron/steel addition acts as coolant because melting is endothermic.

  • The whole process must meet the specification for an individual customer requirement.

  • Dissolved oxygen is removed with aluminium

    • 4Al + 3O₂ ==> 2Al₂O₃ 

  • and then C, Mn and Si etc. can be re–added to a desired specification, plus any other elements, to make a particular steel.

  • Argon (of light bulb fame) is bubbled through to stir the mixture because it so unreactive and most 'stirrers' will melt and dissolve, and change the composition.

  • In the future electric arc furnaces maybe used more to recycle steel. Big carbon electrodes are 'sparked' to melt the scrap iron/steel, lime added to remove impurities as slag. It is possible to use this technology on a small scale to produce 

• Steel is an alloy based on iron.

  • An alloy is a mixture of a metal with at least one other element (metal or non–metal) or compound.

  • The composition of steel, like any other alloy, is crucial in determining its properties.

  • Small differences in composition can have significant effect on the properties of an alloy.

  • Too high a % of C in iron makes it too brittle, but a low % C makes a stronger steel.

  • You need to appreciate the versatile nature of steel by changing its composition and quote some examples.

  • There is a need for excluding impurities eg O, P or S which lead to poor quality material.

  • The common elements added to iron to make steel, apart from carbon,  are usually other transition metals.

• Scrap iron and steel is part of BOS process and is cost effective, recycling reduces costs of (i) ore mining extraction, (ii) possibly overseas transport and (iii) blast furnace reduction of ore. These gains are partly offset by the cost of collecting scrap metal.

  • In the electric arc process only scrap steel is used and is handy technology to produce small batches of particular steel by carefully controlling what scrap goes in.

  • The composition of scrap important, needs to be graded and selected to avoid problems

  • When recycling tin cans, you need to remove the tin and other waste.

    • The cans are shredded and paper/residual food removed, mechanical shredding and magnetic separation can be used,

    • and de–tinning is done by reaction with hot NaOH(aq), after which the valuable tin can be recovered by electrolysis of the 'waste solution'.

  • A particular scrap case study

    • There is a particular need for steel uncontaminated by radioactive isotopes from the nuclear and atomic weapon industries.

    • For this, a useful scrap source is from the German ships sunk at Scapa Flow has proved useful (good geography Q and I don't remember the event!).

The Chemistry of IRON

Iron(II) and Iron(III) chemistry

• The electrode potential chart above highlights the values for various oxidation states of iron.

• The most common oxidation states of iron in its compounds are +2 and +3.

• IRON(II) and IRON(III) Chemistry

• Iron readily dissolves in dilute hydrochloric or sulphuric acid to form iron(II) chloride and iron(II) sulphate respectively. Hydrogen gas is evolved and it is a redox reaction.  

  • The pale green hexaaquairon(II) ion, [Fe(H₂O)₆]²⁺_(aq), is formed.

  • Fe _(s) + 2HCl _(aq) ==> FeCl_{2 (aq)} + H_{2 (g)} 

  • Fe _(s) + H₂SO_{4 (aq)} ==> FeSO_{4 (aq)} + H_{2 (g)} 

  • The redox–ionic equation is: Fe _(s) + 2H⁺ _(aq) ==> Fe²⁺ _(aq) + H_{2 (g)} 

  • Hydrogen ions (H in +1 ox. state) are reduced by electron gain to hydrogen gas (H in 0 ox. state) and iron is oxidised from the 0 ox. state to the +2 ox. state. Note that the lower oxidation state of iron is formed, since neither acid is a strong oxidising agent.

  • The pale green salts FeCl₂.6H₂O and  FeSO₄.7H₂O can be made by careful evaporation and crystallisation of the solution.

  • However, they are readily oxidised by dissolved oxygen to form iron(III) compounds (more on this later).

  • White anhydrous iron(II) chloride can be made by passing hydrogen chloride gas over heated iron.

    •  Fe _(s) + 2HCl _(g) ==> FeCl_{2 (g)} + H_2(g) 

• If chlorine is passed over heated iron, brown anhydrous iron(III) chloride is formed

  • 2Fe _(s) + 3Cl_{2 (g)} ==> 2FeCl_{3 (s)} 

  • An example of 'salt' synthesis by directly combining the constituent elements.

  • Iron(III) chloride is a brown covalently bonded, relatively volatile chloride. Like aluminium chloride, it exists in the solid form as a dimer Fe₂Cl₆, one of the Fe's chlorines acts as a bridge, forming a dative co–ordinate bond with the other iron atom (see diagram below).

  • 

  • Redox reaction: ox. state changes are Fe (0) to (+3), Cl (0) to (–1)

• The iron(III) chloride reacts very exothermically with water to give pungent acrid fumes of hydrogen chloride (anhydrous aluminium chloride is made in the same way and behaves with water in the same way!). Hence the need for dry conditions in their preparation is illustrated below. Its also a very good idea to vent the excess chlorine away safely!

• FeCl_{3 (s)} + 3H₂O _(l) ==> Fe(OH)_{3 (s)} + 3HCl _(g) 

• Some reactions of iron(II) and iron(III) ions:

  • For iron(II) chemistry, a solution of iron(II) sulfate FeSO_4(aq) is suitable for most laboratory experiments investigating the aqueous chemistry of the iron(II) ion.

  • For iron(III) chemistry, a solution of iron(III) chloride FeCl_3(aq) and iron(III) sulfate Fe₂(SO₄)_3(aq) are suitable for most laboratory experiments investigating the aqueous chemistry of the iron(III) ion.

  • The hexaaquairon(II) ion [Fe(H₂O)₆]²⁺_(aq) is pale green.

  • The 'pure' hexaaquairon(III) ion [Fe(H₂O)₆]³⁺_(aq) is pale purple BUT this is NOT usually the main species in aqueous solution.

  • What you normally see is the yellow–light brown–orange coloured complex ion formed from proton transfer to water giving a hydroxo–complex ion (see equation below).

  • This proton transfer process can continue in higher pH media to give the iron(III) hydroxide precipitate (see later) and accounts for why iron(III) salt solutions are acidic.

    • [Fe(H₂O)₆]³⁺_(aq) + H₂O_(l) [Fe(H₂O)₅(OH)]²⁺_(aq) + H₃O⁺_(aq) 

  • With the  alkalis sodium hydroxide or ammonia (no ammine complex formed), iron(II) ions and iron(III) ions produce the respective hydrated hydroxide precipitates. There is no further reaction with excess of either i.e. no complexes formed other than the hydrated hydroxide precipitates. All are acid–base reactions and not redox reactions except that iron(II) compounds can be readily oxidised to iron(III) compounds by the oxygen in air..

    • Fe²⁺_(aq) + 2OH^–_(aq) ==> *Fe(OH)_2(s)   (a precipitation reaction)

      • Iron(II) hydroxide is almost white if oxygen is excluded, but in reality forms up as a 'dirty green' ppt., which on exposure to air rapidly turning brown on oxidation to iron(III) hydroxide by dissolved molecular oxygen.

    • then 4Fe(OH)_2(s)  + O_2(g) + 2H₂O_(l) ==> 4Fe(OH)_3(s)  

      • Fe oxidised (II)==>(III), O reduced (0)==>(–2)

        • Fe(OH)₃ can also be thought of as hydrated iron(III) oxide, Fe₂O₃.xH₂O (x is variable)

    • Fe³⁺_(aq) + 3OH^–_(aq) ==> *Fe(OH)_3(s)

      • Iron(III) hydroxide is an orange–brown ppt ('rust' coloured).

      • *The iron(II) and iron(III) hydroxide complex precipitates can be written as 'complexes' i.e.

        • [Fe(OH)₂(H₂O)₄] or [Fe(OH)₃(H₂O)₃],

      • so the reactions could be written as ligand displacement reactions:

        • [Fe(H₂O)₆]²⁺_(aq) + 2OH^–_(aq) ==> [Fe(H₂O)₄(OH)₂]_(s) + 2H₂O_(l)

        • and [Fe(H₂O)₆]³⁺_(aq) + 3OH^–_(aq) ==> [Fe(H₂O)₃(OH)₃]_(s) + 3H₂O_(l)

        • You can even write intermediate ligand exchange equations, but I've shown the intermediate structures diagrammatically below.

      • NaOH is a strong base, fully ionising to Na⁺ and OH^– ions.

      • NH₃ is a weak base but slightly ionises in water to give sufficient hydroxide ions to give the precipitates.

        • NH_3(aq) + H₂O_(l) NH₄⁺_(aq) + OH^–_(aq)

The sequence of ligand displacement reactions (OH^– for H₂O) which occurs with any alkali e.g. when NaOH_(aq), Na₂CO_3(aq) or NH_3(aq) is added to a solution of an iron(II) salt or iron(III) salt, so, the formation of iron(II) hydroxide and iron(III) hydroxide precipitates are shown pictorially as follows ...

			The sequence of iron(II) and iron (III) hydroxide precipitate formation. Each step is essentially one of proton removal from each complex, from []2+ to []0 and []3+ to []0.

• VIEW ppts. in Appendix 12 with OH^–, NH₃ and CO₃^2–.

  • Aqueous sodium carbonate is weakly alkaline and gives the hydroxide ppts. but excess reagent has no further effect. Again, theses are all acid–base reactions and not redox changes.

  • The iron(II) ion probably gives a mixture of the hydroxide (see above) and carbonate too (a basic carbonate?).

    • Fe²⁺_(aq) + 2OH^–_(aq) ==> Fe(OH)_2(s)

    • Fe²⁺_(aq) + CO₃^2–_(aq) ==> FeCO_3(s)

    • which slowly changes to Fe(OH)₂, which in turn is readily oxidised to Fe(OH)₃ (see above).

  • The iron(III) ion gives the hydroxide and carbon dioxide because the hexa–aqua ion is acidic, (see below and  Appendix 1.).

    • *initially 2[Fe(H₂O)₆]³⁺_(aq) + CO₃^2–_(aq) ==> 2[Fe(H₂O)₅(OH)]²⁺_(aq) + H₂O_(l) + CO_2(g)   

    • and this process of proton donation continues until the [Fe(OH)₃(H₂O)₃]_(s) precipitate is formed

    • No Fe₂(CO₃)₃ is formed because of this acid–base reaction. The acidity of the hydrated iron(III) ion makes it react with the carbonate ion. Note Al³⁺ and Cr³⁺ ions behave in the same way.

  • * The acidity of a the iron hexa–aqua ions can be expressed in a 'Bronsted–Lowry' proton transfer style equation:

    • [Fe(H₂O)₆]ⁿ⁺_(aq) + H₂O_(l) [Fe(H₂O)₅(OH)]^(n–1)+_(aq) + H₃O⁺_(aq)

    • where n = 2 or 3. The overall charge on the complex falls by +1 for each proton transferred as an electrically neutral  water ligand is replaced by a charged hydroxide ion (OH^–) ligand.

    • Water acts as the B–L base (H⁺ acceptor) and the hexa–aqua ion acts as the B–L acid (H⁺ donor) in the deprotonation reaction.

    • When n=3 the 'acid action' is strong enough to react with carbonate ions because of the greater polarising action of the more highly charged Fe³⁺ ion compared to the larger and lower charged Fe²⁺ ion. So, when n=2, the acid donation action is too weak and there is no reaction with carbonates and FeCO₃ can be formed.

• The oxidation of iron(II) ions to iron(III) ions

  • i.e. the reducing action of aqueous iron(II) ions:

    • Fe oxidation state change of +2 to +3

  • (i) Chlorine water readily will oxidise iron(II) to iron(III)

    • 2Fe²⁺_(aq) + Cl_2(aq)  ==> 2Fe³⁺_(aq) +  2Cl^–_(aq)

    • Cl oxidation state change of 0 to –1

    • The pale green of the [Fe(H₂O)₆]²⁺_(aq) ion changes to the orange colour of the [Fe(H₂O)₆]³⁺_(aq)  ion.

    • The chlorine water itself is a very pale green, and changes to the colourless chloride ion, so the colour change associated with the oxidation state change of iron(II) to iron(III) is quite clearly seen.

    • Note that chlorine is a powerful enough oxidising agent to oxidise iron(II) ion to the iron(III) ion, BUT iodine is not a strong enough oxidising agent to achieve this. It is in fact the iron(III) ion that will oxidise the iodide ion, rather than the reverse.

      • The oxidising power series for these two situations is

      • Cl₂ (E^Ø_Cl2/Cl– +1.36V) > Fe³⁺ (E^Ø_Fe3+/Fe2+ +0.77V) > I₂ (E^Ø_I2/I– +0.54V),

      • which of course is numerically paralled by the decreasing values of the standard redox potentials of the half–reactions i.e. becoming less positive as the oxidising power decreases.

      • So, cross–check the reaction the oxidation of iodide ions by iron(III) ions described below.

  • (ii) Iron(II) ions reduce potassium manganate(VII), KMnO₄

    • i.e. the manganate(VII) ion is reduced to the manganese(II) ion

    • (i) MnO₄^–_(aq) + 8H⁺_(aq)  + 5Fe²⁺_(aq) ==> Mn²⁺_(aq) + 5Fe³⁺_(aq) + 4H₂O_(l)

    • The overriding colour change is the bright purple manganate(VII) ion being reduced to a pale colour which is a mixture of the very pale pink manganese(II) ion and the pale orange of the iron(III) ion.

    • more details in manganese(VII) chemistry

  • (iii) Iron(II) ions reduce potassium dichromate(VI), K₂Cr₂O₇

    • i.e. the dichromate(VI) ion is reduced to the chromium(III) ion

    • Cr₂O₇^2–_(aq) + 14H⁺_(aq) + 6Fe²⁺_(aq) ==> 2Cr³⁺_(aq) + 6Fe³⁺_(aq) + 7H₂O_(l)

    • Theoretically, there are actually two simultaneous colour changes.

      • The orange dichromate(VI) ion changes on reduction to the green chromium(III) ion,

      • and the pale green iron(II) ion changes on oxidation to the orange iron(III) ion,

      • so I'm not sure exactly how the colour change you would really observe would pan out!

    • more details in chromium(VI) chemistry

  • Reactions (ii) and (iii) can be used to quantitatively estimate Fe²⁺ ions.

    • See also volumetric redox calculations

• Oxidising action of iron(III) ions:

  • With iodide ions, dark brown solution of iodine (or black solid) formed with iron(II) ions.

    • 2Fe³⁺_(aq) + 2I^–_(aq) ==> 2Fe²⁺_(aq) + I_2(aq/s)

    • This accounts for why iron(III) iodide cannot exist.

    • Oxidation state changes: iron Fe, changes from +3 to +2

    • and iodine I, changes from –1 to 0

    • The orange–brown iron(III) ion becomes the pale green iron(II) ion BUT the latter's colour is obscured by the strong dark colour of the iodine formed.

    • E^Ø for Fe³⁺/Fe²⁺ is +0.77V, E^Ø for I₂/I^– is +0.54V, so Fe³⁺ is a stronger oxidising agent than I₂.

    • Note that chlorine is a powerful enough oxidising agent to oxidise iron(II) ion to the iron(III) ion, BUT iodine is not a strong enough oxidising agent to achieve this. It is in fact the iron(III) ion that will oxidise the iodide ion, rather than the reverse.

      • The oxidising power series for these two situations is

      • Cl₂ (E^Ø_Cl2/Cl– +1.36V) > Fe³⁺ (E^Ø_Fe3+/Fe2+ +0.77V) > I₂ (E^Ø_I2/I– +0.54V),

      • which of course is numerically paralled by the decreasing values of the standard redox potentials of the half–reactions i.e. becoming less positive as the oxidising power decreases.

      • So, cross–check this reaction with the oxidation of iron(II) ions to iron(III) ions by chlorine described above.

  • With zinc, colourless zinc and pale green iron(II) ions are formed. This reaction is usually done in the presence of dil. sulphuric acid.

    • Zn(s) + 2Fe³⁺_(aq) ==> 2Fe²⁺_(aq) + Zn²⁺_(aq)

    • Oxidation state changes: Fe +3 to +2, Zn 0 to +2 (Zn²⁺/Zn is –0.76V, less positive redox potential, so stronger reducing agent than Fe²⁺).

    • The reaction can be used as part of a process to titrate and analyse estimate Fe²⁺ and Fe³⁺ mixtures.

      • See redox titration questions

  • –

• Simple test for aqueous iron(III) ions:

  • Add a few drops of ammonium/potassium thiocyanate solution (NH₄SCN/KSCN).

    • The reaction is NOT given by hexa–aqua iron(II) ions.

  • A blood red cationic complex is formed in a ligand exchange reaction, one ligand is displaced by another.

    • [Fe(H₂O)₆]³⁺_(aq) + SCN^–_(aq) ==> [Fe(H₂O)₅SCN]²⁺_(aq) + H₂O_(l)

    • See also Appendix 9 on colorimetry

  • If fluoride ions (e.g. via KF_(aq)) are added the red colour disappears immediately because a 2nd ligand displacement reaction occurs forming the fluoro–complex ion.

    • [Fe(H₂O)₅SCN]²⁺_(aq) + F^–_(aq) ==> [Fe(H₂O)₅F]²⁺_(aq) + SCN^–_(aq)

    • The fluoride ligand bonds more strongly than the thiocyanate ion.

      • [Fe(H₂O)₆]³⁺_(aq) + SCN^–_(aq) ==> [Fe(H₂O)₅SCN]²⁺_(aq) + H₂O_(l)

        • K_stab = {[Fe(H₂O)₅SCN]²⁺_(aq)} / {[Fe(H₂O)₆]³⁺_(aq)} {SCN^–_(aq)}

        • K_stab = 1.4 x 10² mol^–1dm³ [lg(K_stab) = 2.1]

      • [Fe(H₂O)₆]³⁺_(aq) + F^–_(aq) ==> [Fe(H₂O)₅F]²⁺_(aq) + H₂O_(l)

        • formation of the initial monosubstituted complex ion

        • K_stab = {[Fe(H₂O)₅F]²⁺_(aq)} / {[Fe(H₂O)₆]³⁺_(aq)} {F^–_(aq)}

        • K_stab = 2.4 x 10⁵ mol^–1dm³  [lg(K_stab) = 5.38]

      • K_stab([Fe(H₂O)₅F]²⁺)   >  K_stab([Fe(H₂O)₅SCN]²⁺) so displacement occurs.

      • { } have been used for concentration because [ ] are used in complex ion formulae.

• Iron/iron(III) oxide mixture is used as the main component of the catalyst in the Haber Synthesis of ammonia from nitrogen and hydrogen.

  • N_2(g) + 3H_2(g) ==> 2NH_3(g)

• Some biochemistry of iron

• The biological role of iron complexes haemoglobin, myoglobin and ferritin.

  • Oxygen, O₂, molecules co–ordinate to an iron(II) ion in the haemoglobin (hemoglobin) molecule ('haem' (porphyrin square planar complex diagram to do), which acts as a giant complex ion in transportation systems of the blood. Essential for respiration energy release, ...

    • more details to add

  • Unfortunately carbon monoxide forms a stronger ligand bond than oxygen and will displace it to give CO its well deserved toxic respiration. It only takes a small amount of CO, and a simple ligand exchange reaction to affect the respiratory system!

  • The enzyme catalase is extremely efficient at decomposing hydrogen peroxide molecule in organisms. One proposed mechanism involves a catalytic cycle of iron(III) and iron(IV) complexes e.g. if somewhat simplified ....

  • ENZYME–Fe^III + H₂O₂ ==> ENZYME–Fe^IV=O + H₂O

  • ENZYME–Fe^IV=O + H₂O₂ ==> ENZYME–Fe^III + H₂O + O₂ 

• Other complexes of Fe²⁺ and Fe³⁺ and the cyanoferrate test for iron(ii) and iron(III) ions

  • Iron(II) ions complex with the ethanedioate dicarboxylate anion, a bidentate ligand:

    • [Fe(H₂O)₆]²⁺_(aq) + 2C₂O₄^2–_(aq) ==> [Fe(C₂O₄)₂]^2–_(aq) + 6H₂O_(l)

    • Probably better presented as the compound Fe[Fe(C₂O₄)₂] colour?

  • Iron(III) ions complex with another bidentate ligand, the 1,2–diaminoethane molecule (en)

    • [Fe(H₂O)₆]³⁺_(aq) + 3en_(aq) ==> [Fe(en)₃]³⁺_(aq) + 6H₂O_(l)  colour?

    • K_stab = {[Fe(en)₃]³⁺_(aq)} / {[Fe(H₂O)₆]³⁺_(aq)} [en_(aq)]³

    • K_stab = 3.98 x 10⁹ mol^–3 dm⁹ [lg(K_stab) = 9.6]

  • Both Fe²⁺ and Fe³⁺ ions give octahedral cyano anionic complex ions with cyanide ions.

    • [Fe(H₂O)₆]²⁺_(aq) + 6CN^–_(aq) ==> [Fe(CN)₆]^4–_(aq) + 6H₂O_(l)

      • Fe²⁺ gives the hexacyanoferrate(II) ion  colour?

      • This complex ion can be crystallised as potassium hexacyanoferrate(II), K₄[(CN)₆]

      • If a solution of potassium hexacyanoferrate(III) is added to a solution of iron(II) ions a dark blue precipitate of what was known as Turnbull's Blue is formed.

      • Turnbull's blue is identical to Prussian blue whose formation is described next.

      • Note the use of superscripts II and III to show the oxidation states of the metals in the complex.

      • Turnbull's blue is formed in two stages ...

      • (i) Fe²⁺_(aq) + [Fe^III(CN)₆]^3–_(aq) ==> Fe³⁺_(aq) + [Fe^II(CN)₆]^4–_(aq)

        • The iron atoms exchange oxidation states!

      • (ii) K⁺_(aq) + Fe³⁺_(aq) + [Fe^II(CN)₆]^4–_(aq) ==> K⁺Fe³⁺[Fe^II(CN)₆]^4–_(s)

        • The formation of the dark blue precipitate from potassium hexacyanoferrate(III) can be used as a test for iron(II) ion Fe²⁺_(aq).

    • [Fe(H₂O)₆]³⁺_(aq) + 6CN^–_(aq) ==> [Fe(CN)₆]^3–_(aq) + 6H₂O_(l)

      • Fe³⁺ gives the hexacyanoferrate(III) ion colour?

      • This complex can be crystallised as potassium hexacyanoferrate(III), K₃[(CN)₆]

      • If a solution of potassium hexacyanoferrate(II) is added to a solution of iron(III) ions the dark blue precipitate of Prussian blue is formed.

      • K⁺_(aq) + Fe³⁺_(aq) + [Fe^II(CN)₆]^4–_(aq) ==> K⁺Fe³⁺[Fe^II(CN)₆]^4–_(s)

        • The formation of the dark blue precipitate from potassium hexacyanoferrate(II) can be used as a test for iron(III) ion Fe³⁺_(aq).

        • There is no exchange of the iron oxidation states here.

  • Fe³⁺ ions give another anionic complex in concentrated chloride ion solutions

    • [Fe(H₂O)₆]³⁺_(aq) + 4Cl^–_(aq) ==> [FeCl₄]^–_(aq) + 2H₂O_(l)

    • formation of the tetrahedral shaped tetrachlroroferrate(III) anion. colour?

    • K_stab = {[FeCl₄]^–_(aq)} / {[Fe(H₂O)₆]²⁺_(aq)} [Cl^–_(aq)]⁴

    • K_stab = 8 x 10^–1 mol^–4 dm¹² [lg(K_stab) = –0.097]

  • Both the hexa–aqua ions of iron(II) and iron(III) readily complex with EDTA

    • [Fe(H₂O)₆]²⁺_(aq) + EDTA^4–_(aq) ===> [Fe(EDTA)]^2–_(aq) + 6H₂O_(l)

      • K_stab = {[Fe(EDTA)₃]^2–_(aq)} / {[Fe(H₂O)₆]²⁺_(aq)} [EDTA^4–_(aq)]

      • K_stab = 2.0 x 10¹³ mol^–1 dm³ and lg(K_stab) = 14.3

    • [Fe(H₂O)₆]³⁺_(aq) + EDTA^4–_(aq) ===> [Fe(EDTA)]^–_(aq) + 6H₂O_(l)

      • K_stab = {[Fe(EDTA)₃]^–_(aq)} / {[Fe(H₂O)₆]³⁺_(aq)} [EDTA^4–_(aq)]

      • K_stab = 1.3 x 10²⁵ mol^–1 dm³ and lg(K_stab) = 25.1

    • Note that the more highly charged Fe³⁺_(aq) ion complexes more strongly than the Fe²⁺_(aq) ion.

  • –

• Summary of some complexes–compounds & oxidation states of iron compare 3d–block elements

• –

RUSTING and anti–corrosion chemistry

• The electrochemical processes of RUSTING which is the corrosion of iron to form an iron oxide by an oxidation process which is energetically favourable, and it is the opposite of its extraction by reduction of  iron oxide.

  • The detailed electrochemistry of rusting

    • The half–cell of oxidation of Fe to Fe²⁺ occurs in regions of low oxygen concentration:

      • Fe _(s) – 2e– Fe²⁺ _(aq) (E^Ø = –0.44V)

    • The half–cell reduction of O₂ (+ H₂O + e^–) to OH^– occurs in the oxygen richer regions, via the e^– flow through the iron from the oxidised iron (above): To add diagram

      • O_{2 (aq/g)} + 2H₂O_(l) + 4e^– 4OH^–_(aq)   (E^Ø = +0.44V, in alkali)

      • or O_{2 (aq/g)} + 4H⁺_(aq) + 4e^– 2H₂O_(l)   (E^Ø = +1.23V, in acid)

    • The result is iron(II) hydroxide, which is then oxidised to iron(III) hydroxide or hydrated iron(III) oxide, i.e. orange–brown rust!

      •  Fe²⁺_(aq) + 2OH^–_(aq) ==> Fe(OH)_2(s) (non redox reaction) and then 

        • Fe(OH)_2(s) + O_2(aq/g) ==>  Fe(OH)_3(s) (or Fe₂O₃.xH₂O)

        • Fe(OH)_3(s) + e^– ==>  Fe(OH)_2(s) + OH^–_(aq)   (E^Ø = –0.56V)

• Unfortunately rust flakes off and so it all eventually corrodes away (later xref/contrast ZnO, Al₂O₃, Cr₂O₃ on metal surface, which do not flake away and offer good anti–corrosion properties)

• Factors affecting rate of rusting e.g. the following all speed up the process!

  • decreasing pH, H⁺_(aq) ions remove OH^–_(aq) formed from the reduction of O_2(g/aq), 

  • increased concentration of any ions improves the conductivity of the aqueous media, which is part of 'redox circuit',

  • and if the iron is in contact with a 'less reactive' metal (meaning a more +ve half–cell potential), corrosion rates increase, because the iron is preferentially oxidised with the more –ve half–cell potential.

• Rust protection–inhibition ... examples ... are x–ref with assignment 7 on p174.

  • A plastic or paint physical barrier to exclude water and oxygen (air),

  • Either by (i) dipping in molten zinc, or (ii) electrolysis with Zn²⁺_(aq) solution and the iron/steel object as –ve cathode, galvanising with Zn layer which results in the formation of ZnO layer.

    • The redox chemistry is similar to Fe rusting (see Fig 21) but the layer does not flake away giving a protective layer of zinc oxide.

    • Even if scratched, the Zn with a more –ve half–cell potential is preferentially oxidised.

  • Sacrificial corrosion with blocks of Zn or Mg and relate their 'sacrifice' to their more negative half–cell potentials, i.e. preferentially more favourable oxidation.

    • Fe²⁺_(aq) + 2e^– Fe_(s)   (E^Ø = –0.44V)

    • Zn²⁺_(aq) + 2e^– Zn_(s)   (E^Ø = –0.76V)

    • Mg²⁺_(aq) + 2e^– Mg_(s)   (E^Ø = –2.38V)

    • reminder that the reduction of oxygen to water is a positive redox potential

      • O_{2 (aq/g)} + 2H₂O_(l) + 4e^– 4OH^–_(aq)   (E^Ø = +0.44V, in alkali)

      • or O_{2 (aq/g)} + 4H⁺_(aq) + 4e^– 2H₂O_(l)   (E^Ø = +1.23V, in acid)

    • so all the metal oxidations are feasible BUT the most negative potential will lead to the preferential oxidation i.e. Mg > Zn > Fe.

  • Stainless steel via Cr addition, forms protective layer of chromium(III) oxide.

• History lesson in food preservation: ‘invention’ of the tin can (tin coated steel) ...

  • Tin plating steel offers some corrosion protection of the iron because tin is not a particularly reactive metal (less negative potential).

    • However, early tin cans suffered from preferential oxidation of Fe due to its more –ve potential, through any microscopic defect in the tin layer, or indeed if it got scratched. This was cured by lacquer coating as an extra protective barrier.

    • Fe²⁺_(aq) + 2e^– Fe_(s)   (E^Ø = –0.44V)

    • Sn²⁺_(aq) + 2e^– Sn_(s)   (E^Ø = –0.14V)

  • Still, fruit juice was a problem, carboxylic acids complex with Sn²⁺_(aq) ions, changes Sn_(s)/Sn²⁺_(aq) potential making it more negative than Fe_(s)/Fe²⁺_(aq), so Sn preferentially corrodes, not toxic and contribute to ‘tangy’ taste BUT don’t keep too long as Fe eventually will dissolve too!

  • Complex formation affecting corrosion behaviour. Here tin(II) ions form a complex with carboxylic acids like citric acid (tridentate ligand), by reducing the Sn²⁺_(aq) concentration, the Sn_(s)/Sn²⁺_(aq) half–cell potential is then made more negative that that of iron! so the protective thin layer of tin is sacrificially corrode, then its the iron! Don't worry too much, the rates of reaction are slow, BUT don't keep tinned fruit on the shelf for too long!

• Estimation of iron in iron(II) salts and tablet formulations.

  • The iron(II) compounds are extracted with water, usually iron(II) sulfate, FeSO₄

  • iron(II) ions can be titrated with potassium manganate(VII)

  • worked examples of redox volumetric calculations involving iron analysis

  • more details to add

Scandium * Titanium * Vanadium * Chromium * Manganese * Iron * Cobalt * Nickel * Copper * Zinc * Silver & Platinum

keywords redox reactions ligand substitution displacement balanced equations formula complex ions complexes ligand exchange reactions redox reactions ligands colours oxidation states: iron ions Fe(0) Fe2+ Fe(+2) Fe(II) Fe3+ Fe(+3) Fe(III) FeCl2 FeCl3 FeSO4 Fe2(SO4)3 2Fe  + 3Cl2  ==> 2FeCl3  FeCl3  + 3H2O  ==> Fe(OH)3  + 3 HCl  [Fe(H2O)6]3+ + H2O [Fe(H2O)5(OH)]2+ + H3O+ 4Fe(OH)2 + O2 + 2H2O ==> 4Fe(OH)3 [Fe(OH)2(H2O)4] or [Fe(OH)3(H2O)3] [Fe(H2O)6]2+ + 2OH– ==> [Fe(H2O)4(OH)2] + 2H2O and [Fe(H2O)6]3+ + 3OH– ==> [Fe(H2O)3(OH)3] + 3H2O 2[Fe(H2O)6]3+ + CO32– ==> 2[Fe(H2O)5(OH)]2+ + H2O + CO2  [Fe(OH)3(H2O)3] [Fe(H2O)6]n+ + H2O [Fe(H2O)5(OH)](n–1)+ + H3O+ MnO4– + 8H+ + 5Fe2+ ==> Mn2+ + 5Fe3+ + 4H2O Cr2O72– + 14H+ + 6Fe2+ ==> 2Cr3+ + 6Fe3+ + 7H2O 2Fe3+ + 2I– ==> 2Fe2+ + I2 [Fe(H2O)6]3+ + SCN– ==> [Fe(H2O)5SCN]2+ + H2O [Fe(H2O)5SCN]2+ + F– ==> [Fe(H2O)5F]2+ + SCN– Kstab = {[Fe(H2O)5SCN]2+} / {[Fe(H2O)6]3+} {SCN–}Kstab = {[Fe(H2O)5F]2+} / {[Fe(H2O)6]3+} {F–}Kstab([Fe(H2O)5F]2+)  > Kstab([Fe(H2O)5SCN]2+) [Fe(H2O)6]2+ + 2C2O42– ==> [Fe(C2O4)2]2– + 6H2O Fe[Fe(C2O4)2] [Fe(H2O)6]3+ + 3en ==> [Fe(en)3]3+ + 6H2O Kstab = {[Fe(en)3]3+} / {[Fe(H2O)6]3+} [en]3 [Fe(H2O)6]2+ + 6CN– ==> [Fe(CN)6]4– + 6H2O Fe2+ + [FeIII(CN)6]3– ==> Fe3+ + [FeII(CN)6]4– K+ + Fe3+ + [FeII(CN)6]4– ==> K+Fe3+[FeII(CN)6]4– [Fe(H2O)6]3+ + 6CN– ==> [Fe(CN)6]3– + 6H2O K+ + Fe3+ + [FeII(CN)6]4– ==> K+Fe3+[FeII(CN)6]4– [Fe(H2O)6]3+ + 4Cl– ==> [FeCl4]– + 2H2O Kstab = {[FeCl4]–} / {[Fe(H2O)6]2+} [Cl–]4 [Fe(H2O)6]2+ + EDTA4– ===> [Fe(EDTA)]2– + 6H2O Kstab = {[Fe(EDTA)3]2–} / {[Fe(H2O)6]2+} [EDTA4–] [Fe(H2O)6]3+ + EDTA4– ===> [Fe(EDTA)]– + 6H2O Kstab = {[Fe(EDTA)3]–} / {[Fe(H2O)6]3+} [EDTA4–]  oxidation states of iron, redox reactions of iron, ligand substitution displacement reactions of iron, balanced equations of iron chemistry, formula of iron complex ions, shapes colours of iron complexes  Na2CO3 NaOH NH3

Advanced Level Inorganic Chemistry of Iron – A level Revision notes to help revise for GCE Advanced Subsidiary Level AS Advanced Level A2 IB Revise AQA GCE Advanced Level Chemistry OCR GCE Advanced Level Chemistry Edexcel GCE Advanced Level Chemistry Salters AS A2 Chemistry CIE Chemistry, WJEC GCE AS A2 Chemistry, CCEA/CEA GCE AS A2 Chemistry revising courses for pre–university students (equal to US grade 11 and grade 12 and AP Honours/honors level courses)

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Alphabetical Index for Science Pages Content A B C D E F G H I J K L M N O P Q R S T U V W X Y Z

Scandium * Titanium * Vanadium * Chromium * Manganese * Iron * Cobalt * Nickel * Copper * Zinc * Silver & Platinum

Introduction 3d–block Transition Metals * Appendix 1. Hydrated salts, acidity of hexa–aqua ions * Appendix 2. Complexes & ligands * Appendix 3. Complexes and isomerism * Appendix 4. Electron configuration & colour theory * Appendix 5. Redox equations, feasibility, E^ø * Appendix 6. Catalysis * Appendix 7. Redox equations * Appendix 8. Stability Constants and entropy changes * Appendix 9. Colorimetric analysis and complex ion formula * Appendix 10 3d block – extended data * Appendix 11 Some 3d–block compounds, complexes, oxidation states & electrode potentials * Appendix 12 Hydroxide complex precipitate 'pictures', formulae and equations