(2) Data Tables for the physical and chemical properties of iron - a 3d block transition element

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

Elect. pot. = standard electrode potential data for iron (E^Ø at 298K/25^oC, 101kPa/1 atm.)

na = data not applicable or not available for iron (less common oxidation state of iron)

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 steels for making electromagnets.

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

  • 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.

  • The haemoglobin molecule, an animal's oxygen carrier, contains four protein complexes containing an Fe²⁺ ion. In the initial binding of an oxygen molecule to the Fe²⁺ complex ion in haemoglobin it is the iron(II) ion - induced dipole force which brings the O₂ molecule towards the Fe²⁺ ion.

  • Iron deficiency causes anaemia, and sometimes you can take iron tablets to remedy this condition.

  • Plants require iron for the synthesis of chlorophyll.

  • See also (11) More on catalysis by iron or iron compounds and more on the biological role of iron

• Examples of catalysis

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

    • N_2(g) + 3H_2(g) === Fe/Fe₂O_{3 catalyst} ===> 2NH_3(g)

  • The iron(II) and iron(III) ion catalyse the oxidation of iodide ions by peroxodisulfate ions is described under homogeneous catalysis in Appendix 6. It is an example of homogeneous catalysis to reduce the activation energy by providing an alternative pathway.

  • See also (11) More on catalysis by iron or iron compounds and more on the biological role of iron

• Moss killer

  • Ammonium iron(II) sulfate (commercially called by its old name of ferrous ammonium sulfate) is a crystalline double salt made from iron(II) sulfate FeSO₄ and ammonium sulfate (NH₄)₂SO₄

    • from which you can make the anhydrous form Fe(NH₄)₂(SO₄)₂, but can be crystallised as ammonium iron(II) sulfate hexahydrate, Fe(NH₄)₂(SO₄)₂.6H₂O, known as Mohr's salt which contains the hexaaquairon(II) complex ion. You should able to see the +/- ion charge balance in the formulae.

  • This double salt dehydrates moss, and, because moss rapidly absorbs nutrients through its leaves (rather than its roots), the moss also suffers from too much iron, killing it.

  • In comparison, the grass benefits from the ammonium sulfate providing it with the essential of nitrogen, sulfur and iron, and iron stimulates chlorophyll production.

• BUT for advanced A level, a few details may be required.

  • Reminders of the sort of reactions that can go on in a blast furnace powered by coke (mainly carbon)

  • Fe₂O₃(s) + 3C(g) ==> 2Fe(l/s) + 3CO(g)

  • 2Fe₂O₃(s) + 3C(g) ==> 4Fe(l/s) + 3CO₂(g)

  • Fe₂O₃(s) + 3CO(g) ==> 2Fe(l/s) + 3CO₂(g)

  • Fe₃O₄(s) + 4C(s) ==> 3Fe(l/s) + 4CO(g)

  • Fe₃O₄(s) + 2C(s) ==> 3Fe(l/s) + 2CO₂(g)

  • haematite iron ore is iron(III) oxide, Fe₂O₃ (+3 oxidation state)

    • in the ore reduction reactions the redox changes are

    •  Fe²⁺/Fe³⁺ (+2/+3) ==>Fe (0);  C (0) ==> CO (C, +2) or C (0) ==> CO₂ (C, +4)

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

    • note that Fe₃O₄ = 2Fe³⁺ + Fe²⁺ + 4O^2–

    • so one iron(II) ion at +2 and two iron(III) ions at +3  and O is -2 in terms of oxidation states.

• 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.

  • Sulfur is removed early in the process using magnesium:

    • Mg + S ==> MgS (the magnesium sulfide 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 e.g. 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 in the post-WWII period.

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

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

The electrode potentials involving iron ions correspond to hydrated complex ions where the ligands are water, oxide or hydroxide.

As you can see from the chart, changing either the ligand or the oxidation state, will also change the electrode potential for that half-reaction involving a iron ion.

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

Electron configuration of Fe²⁺ is [Ar]3d⁶

Electron configuration of Fe³⁺ is [Ar]3d⁵

PLEASE note: The electrode potentials (E^Ø) for iron chemistry in the text sometimes vary from the chart above - apologies, but data sources for iron chemistry can vary! Any discrepancy shouldn't significantly affect any electrode potential calculation outcomes e.g. feasibility (email if concerned?).

(6) Some reactions of iron metal

• Iron readily dissolves in dilute hydrochloric or sulfuric acid to form iron(II) chloride and iron(II) sulfate 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 oxidation state) are reduced by electron gain to hydrogen gas (H in 0 oxidation state) and iron is oxidised from the 0 oxidation state to the +2 oxidation 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. It exists in the solid form as a covalent molecular lattice of the dimer Fe₂Cl₆.

    • So, strictly speaking the equation should be written as: 2Fe _(s) + 3Cl_{2 (g)} ===> Fe₂Cl_6(s) 

    • 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: oxidation 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 as illustrated above. Its also a very good idea to vent the excess chlorine away safely too!

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

    • If the iron - halogen experiment is repeated with bromine the reaction is less vigorous, but iron(III) bromide is still formed.

      • The exothermic nature of the reaction may or may not be seen?

      • The dimer molecules are present in the brown solid.

      • 2Fe(s) + 3Br₂(g) ===> Fe₂Br₆(s)

      • The reaction is easily demonstrated by warming a little bromine with iron wool in a fume cupboard!

  • When iron wool is heated with iodine there is little reaction, a small amount of iron(II) iodide is formed.

    • Fe(s) + I₂(s) ===> FeI₂(s)

    • Fe³⁺ is sufficient in oxidising power to oxidise an iodide ion to iodine, so FeI₂ is formed, not FeI₃.

    • This is theoretically predictable from the half-cell electrode potentials.

    • E^Ø_I2/I– +0.54V  <  E^Ø_Fe2+/Fe3+ +0.77V.

    • The more positive Fe²⁺/Fe³⁺ potential will oxidise iodide to iodine.

    • If by chance any FeI₃ was formed, it would readily decompose to FeI₂ and I₂.

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

• Formation of a complex ion

• Above is an old fashioned valence-bond approach to describing the bonding in a transition metal complex.

• For iron(II) aqueous 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) aqueous 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.

• It has an octahedral shape and a co-ordination number of 6 from 6 unidentate ligands.

• 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)

  • This is why iron(III) salt solutions are weakly acidic (typically pH 2-6) and more acidic than iron(II) salt solutions because the Fe³⁺ ion has a higher charge density to facilitate proton transfer.

  • The Fe³⁺ ion has a greater charge and smaller radius than the Fe²⁺ ion.

  • See Appendix 1. Hydrated salts, acidity of hexa–aqua ions and patterns explained

• 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 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 (effervescence observed) because the hexa–aqua Fe³⁺ 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.

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

Useful half-cell electrode potential data to help understand the redox chemistry of iron

• 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 hexaaquairon(III) ion is actually a very pale violet, BUT, it acts as a weak acid and protonates water to form complexes such as [Fe(H₂O)₅OH]³⁺_(aq), the pentaaquamonohydroxoiron(III) ion which give orange to brown colours - think also of the brown hydrated precipitate complex of iron(III) hydroxide.

  • 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 paralleled 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 and iron(II) oxidised to iron(III);

  • 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 and iron(II) oxidised to iron(III);

  • 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 (Questions and worked answers)

• (iv) Oxidation of iron(II) to iron(III) by hydrogen peroxide solution

  • Mixing iron(II) salt solutions with acidified hydrogen peroxide will effect the oxidation

  • 2Fe²⁺(aq)  +  H₂O₂(aq)  +  2H⁺(aq)  ===>  2Fe³⁺(aq)  +  2H₂O(l)

  • Hydrogen peroxide solution is a more powerful oxidising agent than the iron(III) ion.

(9) Oxidising action of iron(III) ions:

Useful half-cell electrode potential data to help understand the redox chemistry of iron

• (i) 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 in solution or precipitate.

  • 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 paralleled 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.

• (ii) With zinc, colourless zinc ions and pale green iron(II) ions are formed

  • This reduction reaction is usually done in the presence of dilute sulfuric 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 (Questions and worked answers)

• (iii) Oxidation of thiosulfate

  • Iron(III) ions oxidise the thiosulfate ion to the tetrathionate ion (both anions are colourless), but apparently a dark-violet co-ordination complex is formed initially.

  • 2Fe³⁺(aq)  +  2S₂O₃^2-(aq)  ===>  2Fe²⁺(aq)  +  S₄O₆^2-(aq)

(10) Simple tests for aqueous iron(II) and iron(III) ions:

All the equations are in section (7)

 Iron(II) ions give a dark green precipitate with dilute sodium hydroxide solution - turns brown in air - oxidation to iron(III) hydroxide.

 Iron(III) ions give a brown precipitate with dilute sodium hydroxide solution.

The hydroxide precipitate colours can easily distinguish between an aqueous solutions containing Fe²⁺ or Fe³⁺, but this is a specific test for the aqueous iron(III) ion, assuming the compound is soluble in water.

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

  • The reaction is NOT given by hexaaquairon(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)

  • This is an example of a ligand substitution (ligand exchange) reaction.

  • The thiocyanate ligand displaces a water ligand, no change in shape, oxidation state or coordination number.

  • Both iron(III) complex ions are octahedral in shape with a co-ordination number of 6.

  • The oxidation state of iron remains at +3, but the overall charge of the iron thiocyanato complex formed drops to 2+ because a + is cancelled out by the - of the thiocyanate ion.

  • In most ligand exchange reactions there is no change in oxidation state unless a reducing agent or oxidising agent is present.

  • 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.

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

  • The fluoride ligand bonds more strongly than the thiocyanate ion, and comparing the reactions ...

    • (ii)  [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 = 1.40 x 10² mol^–1dm³ [lg(K_stab) = 2.1]

      • Remember [H₂O] is not included in the equilibrium expression.

    • (iii) [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 = 2.40 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.

    • All the three iron complex ions above are octahedral in shape with a co-ordination number of 6.

    • You can calculate the equilibrium constant for reaction (i) from the K_stab values for reactions (ii) and (iii).

    • K_equilib  for (i) is derived by dividing  K_stab (ii) by K_stab (iii)

    • K_equilib  = ([[Fe(H₂O)₅F]²⁺_(aq)] x [SCN^–_(aq)]) / ([[Fe(H₂O)₅SCN]²⁺_(aq)] x [F^–_(aq)]) =  2.4 x 10⁵ / 1.4 x 10² = 1.71 x 10³ (no units)

(11) More on catalysis by iron or iron compounds and more on the biological role of iron

See also Appendix 6. Catalysis - heterogeneous and homogeneous transition metal examples explained

and kinetics section 5.2 Heterogeneous and homogeneous catalysis reaction mechanisms

(i) An 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)

The iron/iron oxide provide 'active sites' to break open the strong covalent bonds of the hydrogen (H–H) and nitrogen (N≡N) molecules.

Remember that a catalyst increases the reaction rate, but cannot change the position of an equilibrium, so no increase in % product yield.

(ii) The decomposition of hydrogen peroxide by iron(II) or iron(III) salts

If you add any iron(II) or iron(III) sulfate solution to a hydrogen peroxide solution, effervescence  is seen as oxygen is given off after the hydrogen peroxide decomposes.

2H₂O₂(aq)  ===>  2H₂O(l)  +  O₂(g)

The reaction involves a catalytic redox recycling of Fe²⁺ Fe³⁺ 'complexes', so it doesn't matter which salt you start with!

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 (haemoglobin) molecule.

  • The molecule is a protein-porphyrin ring complex (haem)

  • The haemoglobin (haem) molecule acts as a multi/polydentate ligand with iron(II) ions in blood chemistry.

  • In an extremely simplified form the structure is an iron(II) complex: [protein–Fe^II–O₂].

  •  ('haem' (includes a porphyrin square planar structure), acts as a giant complex ion in transportation systems of the blood i.e. the transfer of oxygen from the lungs to all cells of the body. Transportation of oxygen is essential for respiration - energy release to power the biochemistry of most living organisms.

    • See also The uv-visible absorption spectrum of haemoglobin (porphyrin pigment)

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

    • [protein–Fe^II–O₂]  +  CO  ===>  [protein–Fe^II–CO]  +  O₂

• 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 

• Iron complexes in the cytochrome system are essential for human metabolism. They act as electron transfer systems in complex redox reactions that power cellular respiration (energy production) and the detoxification of harmful substances.

(12) Other complexes of Fe²⁺ and Fe³⁺ K_stab values and the cyanoferrate test for iron(II) and iron(III) ions

• Iron(II) ions complex with the ethanedioate dicarboxylate anion (oxalate ion), a bidentate ligand (shown on the right):

  • [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? K_stab?

• Iron(III) ions complex with another bidentate ligand, the 1,2–diaminoethane molecule (H₂NCH₂CH₂NH₂ = en, very short for the old name of ethylenediamine)

  • [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 yellow coloured hexacyanoferrate(II) ion

    • 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 (brown-deep red colour)

    • The two iron(III) complexes are octahedral in shape with a co-ordination number of 6.

    • The overall electrical charge on the iron complex changes from 3+ to 3- (from +3 and 6 x -1).

    • This is an example of an iron complex ligand exchange reaction where six cyanide ions replace six water molecules.

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

    • The bonding details of this iron(III) complex are shown below.

    • This is an old fashioned valence-bond approach to describing the bonding in a transition metal complex.

    • 

    • 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, which turn's out to be the same as Turnbull's Blue .

    • 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.

      • Note the use of superscript II to indicate the different oxidation states of iron.

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

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

    • In this ligand exchange reaction, the iron(III) complex ion shape changes from octahedral to tetrahedral, the co-ordination number changes from 6 to 4, but the oxidation state of iron remains at +3.

    • The overall electrical charge on the chloro complex is -1 (from 3+ and 4 x -1).

    • Its likely that the more bulky chloride ion (radius Cl > O) 'forces' the formation of the tetrahedral shape of this iron complex ion, rather than a square planar shaped complexes.

    • 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 (hexadentate ligand)

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

      • K_stab = [[Fe(EDTA)]^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)]^–_(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 and smaller radius Fe³⁺_(aq) ion complexes more strongly than the Fe²⁺_(aq) ion (Fe³⁺ has a greater charge density).

      • This complex in used in iron deficiency tablets to combat anaemia.

  • The reactions with multi-dentate ligands like diamines, oxalate and EDTA are referred to as chelation reactions (sequestration is a term also used), where  the dative bonds 'enclose' the central metal ion.

  • They often readily displace mono-dentate ligands like H₂O and NH₃, partly because the displaced mono-dentate ligands causes a great increase in entropy.

  • See Appendix 8. Stability Constants and entropy changes

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

• 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, anode reaction)

    • 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):

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

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

      • Either way, the electrode potential predicts that iron can be oxidised by aqueous oxygen.

    • 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) (not balanced)

      • Fe(OH)_3(s) + e^– ===>  Fe(OH)_2(s) + OH^–_(aq)   (half-reaction, 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),

  • Or 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 above) 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 (data below).

  • 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 to steel alloys, 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 simply 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!

(14) The volumetric estimation of iron e.g. in iron(II) salts and tablet formulations.

• (i) The iron(II) compounds are extracted with water from various formulations, usually iron(II) sulfate, FeSO₄ e.g. in iron deficiency medication tablets or moss killer.

• (ii) Samples of steel can be dissolved in dilute sulfuric acid and the solution titrated with potassium manganate(VII) solution to determine the iron content.

• You can also titrate iron(II) ions with standardised potassium dichromate using a redox indicator.

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

• See section (16) Volumetric titration questions involving the analysis for iron (with answers)

• worked examples of general redox volumetric calculations

(15) The vertical connection of iron with the other d-block elements of Group 8 (IUPAC designation)

You would expect some similarities in the chemistry of iron, ruthenium and osmium which have the same outer electron configuration as iron.

Which bar ruthenium, have an identical outer electron configuration of nd⁶s² or nd⁷s¹ (n = 3 to 6)

(16) Volumetric titration questions involving the analysis for iron

Calculation examples adapted and re-edited from worked examples of redox volumetric calculations

Question 1: Given the following two half–reactions: (Q1 can be done as an experimental 'word–fill' version)

need to add method Q

Question 1 has many parts covering the titration of iron(II) ions with a standard solution of potassium manganate(VII) and the problems are solved.

The relative atomic mass of iron = 55.9

Q1(e) was a late addition and the worked out answers are better presented than maybe others on this page?

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

and (ii) Fe³⁺_(aq) +  e^– ==> Fe²⁺_(aq)

(a) Construct the fully balanced redox ionic equation for the manganate(VII) ion oxidising the iron(II) ion

(b) 24.3 cm³ of 0.0200 mol dm^–3 KMnO₄ reacted with 20.0 cm³ of an iron(II) solution acidified with dilute sulfuric acid.

(i) Calculate the molarity of the iron(II) ion.

hint: first work out mol of manganate(VII), then from equation mol of iron(II), then molarity

(ii) How do recognise the end–point in the titration?

(c) Calculate the percentage of iron in a sample of steel wire if 1.51 g of the wire was dissolved in excess of dilute sulfuric acid and the solution made up to 250 cm³ in a standard graduated flask.

A 25.0 cm³ aliquot of this solution was pipetted into a conical flask and needed 25.45 cm³ of  O.0200 mol dm^–3 KMnO₄ for complete oxidation.

hints as above, then mass of iron, but need to take into account the dilution.

(d) Suggest reasons why the presence of dilute sulfuric acid is essential for an accurate titration and why dilute. hydrochloric and nitric acids are unsuitable to be used in the context of this titration.

(e) Calculation based on the analysis of a soluble iron(II) salt

8.25g of an iron(II) salt was dissolved in 250 cm³ of pure water. 25.0 cm³ aliquots were pipetted from this stock solution and titrated with 0.0200 mol dm^–3 potassium manganate(VII) solution.

The titration values obtained were 23.95 cm³, 23.80 cm³ and 23.85 cm³.

(i) What titration value should be used in the calculation and why?

(ii) Calculate the moles of manganate(VII) used in the titration.

(iii) calculate the moles of iron(II) ion titrated

(iv) Calculate the mass of iron(II) titrated

(v) Calculate the total mass of iron in the original sample of the iron(II) salt.

(vi) calculate the % iron in the salt.

Question 9: 13.2 g of iron(III) alum were dissolved in water and reduced to an iron(II) ion solution by zinc and dilute sulfuric acid. The mixture was filtered and the filtrate and washings made up to 500 cm³ in a standard volumetric flask.

If 20.0 cm³ of this solution required 26.5  cm³ of 0.0100 mol dm^–3 KMnO₄ for oxidation.

(a) give the ionic equation for the reduction of iron(III) ions by zinc metal.

(b) Calculate the percentage by mass of iron in iron alum.

hints: See Q1, particularly Q1(e)

ANSWERS and WORKING

Question 11: Calculate x in the formula FeSO₄.xH₂O from the following data:

12.18 g of iron(II) sulfate crystals were made up to 500 cm³ acidified with sulfuric acid.

25.0 cm³ of this solution required 43.85 cm³ of 0.0100 mol dm^–3 KMnO₄ for complete oxidation.

hints: See Q5, but manganate(VII) instead of dichromate(VI) and different ratio.

ANSWERS and WORKING

Question 13: 2.68 g of the salt iron(II) ethanedioate, FeC₂O₄, were made up to 500 cm³ of acidified aqueous solution. 25.0 cm³ of this solution reacted completely with 28.0 cm³ of 0.0200 mol dm^–3 potassium manganate(VII) solution.

Calculate the mole ratio of KMnO₄ to FeC₂O₄ taking part in this reaction.

Give the full redox ionic equation for the reaction - tricky to balance - triple checks of - oxidation numbers up/down, ionic charge (= both sides of equation), and atom count.

hints: calculate molar mass of salt and number of moles, calculate mol of manganate(VII) in titration.

ANSWERS and WORKING

Question 15: Calculate the molarities of iron(II) and iron(III) ions in a mixed solution from the following data.

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

(i) 25.0 cm³ of the original mixture was acidified with dilute sulfuric acid and required 22.5 cm³ of 0.0200 mol dm^–3 KMnO₄ for complete oxidation.

hints: calculate moles of Fe(II) then molarity from given volume, Fe(III) NOT titrated

(ii) a further 25.0 cm³ of the original iron(II)/iron(III) mixture was reduced with zinc and acid and it then required 37.6 cm³ of the KMnO₄ for complete oxidation.

hints: calculate moles of Fe(II) + Fe(III), then total molarity from given volume, all Fe titrated, a little subtraction in the end.

ANSWERS and WORKING

Question 16: A piece of rusted iron was analysed to find out how much of the iron had been oxidised to rust [hydrated iron(III) oxide].

A small sample of the rusted iron was dissolved in excess dilute sulfuric acid to give 250 cm³ of solution.

The solution contains Fe²⁺ ions from the unrusted iron dissolving in the acid, and, Fe³⁺ ions from the rusted iron.

(a) 25.0 cm³ of this solution required 16.9 cm³ of 0.0200 mol dm^–3 KMnO₄ for complete oxidation of the Fe²⁺ ions.

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

Calculate the moles of Fe²⁺ ions in the sample titrated.

(b) To a second 25.00 cm³ of the rusted iron solution an oxidising agent was added to convert all the Fe²⁺ ions present to Fe³⁺ ions.

The Fe³⁺ ions were titrated with a solution of EDTA^4–_(aq) ions and 17.6 cm³ of a 0.100 mol dm^–3 EDTA were required.

Assuming 1 mole of EDTA reacts with 1 mole of Fe³⁺ ions, calculate the moles of Fe³⁺ ions in the sample.

(c) From your calculations in (a) and (b) calculate the ratio of rusted iron to unrusted iron and hence the percentage of iron that had rusted.

ANSWERS and WORKING

Question 20 Lawn sand containing the salt iron(II) sulfate is used to treat moss.

2.50 g of the lawn sand was mixed with dilute sulfuric acid to extract the iron(II) salt.

In a volumetric titration, it required 24.50 cm³ of a 0.0200 mol dm^-3 solution of potassium manganate(VII) to fully oxidise the iron salt.

(a) Write out the equation for the redox reaction involved in the titration.

(b) Calculate the % by mass of Fe²⁺ ions in the lawn sand. Atomic mass of iron = 55.8

Hint: First calculate mol of KMnO₄ used in titration, then from redox equation calculate the mol of Fe²⁺ titrated ...

ANSWERS and WORKING

Question 21 Analysing a medicinal iron tablet containing a soluble iron(II) salt.

(For people showing signs of iron deficiency e.g. low red blood cell count, suffering from anaemia).

(a) Given a sample of 'iron' tablets, describe how you would analyse them to find the average mass of iron in each tablet.

(b) make a titre 24.50

A 1.02 g iron tablet was dissolved in dilute sulfuric acid and titrated with a standard solution of potassium dichromate(VI).

With a suitable redox indicator, it took 23.80 cm³ of a 0.02 mol dm^-3 solution of the dichromate reagent to completely oxidise all the Fe²⁺ ions in the tablet.

(a) Give the full redox equation of the titration.

(b) Calculate the % by mass of Fe²⁺ ions in the tablet. Atomic mass of iron = 55.8

Hint: First calculate mol of K₂Cr₂O₇ used in titration, then from redox equation calculate the mol of Fe²⁺ titrated ...

ANSWERS and WORKING