Transition metals: Some chemistry of iron(1)

Some chemistry of Iron(1):

Edexcel

24. be able to record observations and write suitable equations for the reactions of Cr³⁺(aq), Fe²⁺(aq), Fe³⁺(aq), Co²⁺(aq) and Cu²⁺(aq) with aqueous sodium hydroxide and aqueous ammonia, including in excess.

25. be able to write ionic equations to show the difference between ligand exchange and amphoteric behaviour for the reactions in (24) above.

34. understand the role of Fe²⁺ ions in catalysing the reaction between I− and S₂O₈^2— ions.

Manganate(VII) with iron (II) titration self-indicating

catalyst in Haber process

AQA

Exchange of the ligand H₂O by Cl– can involve a change of co-ordination number (e.g. Fe³⁺(aq), Co²⁺(aq) and Cu²⁺(aq).

Haem is an iron(II) complex with a multidentate ligand.

Oxygen forms a co-ordinate bond to Fe(II) in haemoglobin, enabling oxygen to be transported in the blood.

Carbon monoxide is toxic because it replaces oxygen co-ordinately bonded to Fe(II) in haemoglobin.

The redox titrations of Fe²⁺(aq) and C₂O₄^2– with MnO₄^—

Students should be able to perform calculations for these titrations and similar redox reactions.

Examples include, finding:

•   the mass of iron in an iron tablet 


•   the percentage of iron in steel 


•   the Mr of hydrated ammonium iron(II) sulfate 


•       Fe is used as a heterogeneous catalyst in the Haber process.

In aqueous solution, the following metal-aqua ions are formed:

[M(H₂O)₆]²⁺, limited to M = Fe and Cu

[M(H₂O)₆]³⁺, limited to M = Al and Fe

The acidity of [M(H₂O)₆]³⁺ is greater than that of [M(H₂O)₆]²⁺

Students should be able to:

•   explain, in terms of the charge/size ratio of the metal ion, why the acidity of [M(H₂O)₆]³⁺ is greater than that of [M(H₂O)₆]²⁺ 


•   describe and explain the simple test-tube reactions of M²⁺(aq) ions, limited to M = Fe and Cu, and of M³⁺(aq) ions, limited to M = Al and Fe, with the bases OH–, NH3 and CO₃^2— 


•       Students could carry out test-tube reactions of metal-aqua ions with NaOH, NH₃ and Na₂CO₃

Halogen carrier in arene organic chemistry

OCR

Redox reactions


(k) redox reactions and accompanying colour changes for:

(i) interconversions between Fe2+ and Fe3+

Fe2+ can be oxidised with H+/MnO4– and Fe3+ reduced with I–

Redox chemistry of iron and its compounds

1. Iron production in the Blast Furnace

Iron production is massive across the world.  It is still the most important piece of chemistry happening.  Nations across the world rely on iron and steel as construction materials for buildings and vehicles. 

And the material starts as iron(III) oxide Fe₂O₃

Fe₂O₃    +   3CO   ⟶     3CO₂   +   2Fe

Carbon monoxide generated in a blast furnace from the action of coke in oxygen–enriched air reduces the iron(III)oxide to iron which is then taken for processing into steel. 

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2. Relative reactivity of iron and copper

One of the most useful characteristics of iron is that it is not highly reactive and can be easily handled and stored as a solid unlike many more reactive metals like sodium which require protection on storage and handling. 

Furthermore, this low reactivity means that it does not corrode quickly and its corrosion can be easily prevented in several ways e.g. by painting or oiling or galvanising or using sacrificial protection.

Iron is easily plated if it is placed in a solution of copper ions. 

The redox chemistry is described in the equation below. 

CuSO₄   +    Fe     ⟶       FeSO₄     +   Cu

The reactivity series results from this difference in ability to react.

Highly reactive metals good at losing electrons from their outer shells we put top of the list and those metals unwilling to lose electrons are at the bottom. 

In the chart below, reactivity series is related to electrode potential.

3. Thermit reaction

This is the chemistry used to weld rail–track into continuous rail in situ.

2Al   +    Fe₂O₃      ⟶       2Fe     +     Al₂O₃

Aluminium powder reduces iron oxide to iron is the basic reaction but of course the metal mix has to match the steel qualities of the rails or else the weld would crack under use. 

5. Precipitation Reactions of iron(II) and iron(III).

With aqueous sodium hydroxide:  NaOH(aq)

Addition of sodium hydroxide solution to a solution of iron(I) or iron (III) ions produces distinctive precipitates both of which are insoluble in excess sodium hydroxide solution. 

Fe²⁺       +   2OH^—       ⟶         Fe(OH)₂(s)

                                                Pale green

The iron(II) hydroxide precipitate is a very pale green at first which turns darker green on standing due, actually, to oxidation if the iron(II) ions by dissolved oxygen in the water.  You can see in the photo below that the iron(II) hydroxide has darkened near the surface of the solution due to action of atmospheric oxygen.

Fe³⁺       +   3OH^—       ⟶         Fe(OH)₃(s)

                                                Dark brown

The iron(III) hydroxide precipitate is a dark rust brown in colour and it remains so on standing. 

These different colours serve to distinguish iron(II) from iron(III) ions. 

These precipitates are pictured below:

Similar reactions occur with alkaline aqueous ammonia (NH₃(aq))

6. Acidity of iron(III) ions:

The acidity of [Fe(H₂O)₆]³⁺ is greater than that of [Fe(H₂O)₆]²⁺

Iron(III) ions in aqueous solution are excellent proton donors with a significant pKa value. 

[Fe(H₂O)₆]³⁺       ⟶        [Fe(H₂O)₅ OH]²⁺      +      H⁺

pale lilac                               orange

Iron(II) ions by contrast are nowhere near as effective as proton donors and have a significantly different and larger pKa value. 

How can we explain the difference in acidity?

The difference in pKa values is down to the difference in charge/ size ratios or charge density values.

The iron(III) ion has a smaller radius and a greater charge giving it a greater charge density.

The greater the charge density (charge/size ratio) means that the iron(III) ion has greater polarising power acting on the water molecule ligands.

The effect is to distort the electron distribution in the water molecules and to therefore to weaken the OH bond in the water molecules and make the loss of a proton more likely. 

The diagram below shows how this works with large sodium and small iron(III) ions.

My next post will complete this brief survey of the chemistry of iron and some of its compounds.