The Steel Story

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Rusting 2

OXIDATION: Fe(s) --> Fe2+(aq)+ 2e-

REDUCTION: ½ O2(g) + H2O(l) + 2e- --> 2OH-(aq)

Corrosion is always greater at the centre of a water drop or under a layer of paint because oxygen supply is limited. Rust forms in a secondary process as Fe2+ and OH- ions diffuse away from the metal. 

Fe2+(aq)+ 2OH-(aq) --> Fe(OH)2(s)

Fe(OH)2(s) ----------> Fe2O3 . xH2O(s)


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Preventing Corrosion

  • Barrier Protection- prevents oxygen and/ or water coming into contact with iron and steel, eg. painting, greasing and oiling.
  • Galvanising- Covering in a thin layer of zinc that oxidises. 
  • Sacrificial Protection- attaching blocks of a more reactive metal (eg. zinc) to large iron structures (eg. ships). An electrochemical cell is formed and the reactive metal corrodes as a preference to the iron. 
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The d block- Transition Metals


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The d block- Transition Metals 2

Sc [Ar] 3d14s2 

Ti [Ar] 3d24s2 

V [Ar] 3d34s2

Cr [Ar] 3d54s1 

 Mn [Ar] 3d54s2 

Fe [Ar] 3d64s2

Co [Ar] 3d74s2

Ni [Ar] 3d84s2

Cu [Ar] 3d104s1

Zn  [Ar] 3d104s2

A transition metal is one which forms one or more stable ions which have incompletely filled d orbitals.

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The d block- Transition Metals 3

Physical Properties

  • good conductors of heat and electricity
  • dense
  • high melting and boiling points
  • hard
  • durable
  • high tensile strength

Chemical Properties

  • variable oxidation states
  • catalytic activity
  • formation of complexes
  • formation of coloured compounds
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Transition Metals- Variable Oxidation States (Redo

Transition metal have a range of oxidation states. This is because the differences between successive ionisation enthalpies in the 3d and 4s subshells are relatively small, so multiple electron loss is possible. 

In lower oxidation states, the elements exist as simple ions eg (Cu2+, Fe2+, Fe3+). In the higher oxidation states they are covalently bonded to electronegative elements such as O and F, forming anions, (eg Cr2O72-, MnO4-). 

In REDOX- Compounds containing transition metals in high oxidation states tend to be oxidising agents and vice versa. 

 Potassium Manganate (VII)Titrations

  • Strong oxidising agent,used to find concentrations of solutions containing Fe2+ ions or hydrogen peroxide. No indicator needed as deep purple colour disappears as it reacts.

Iodine- Thiosulfate Reactions

  • Near the end point, starch solution is added which gives an intense blue/ black colour which disappears at the end point. 
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Transition Metals- Variable Oxidation States (Redo

Vanadium Oxidation States and Colours. Your Brother Grows Viagra. 


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Transition Metals- Catalytic Activity

Transition metals act as good catalysts as they can change their oxidation state and adsorb substances onto their surface and activate them. 

Transition METALS and their COMPOUNDS act as heterogeneous catalysts, providing a surface for gaseous reactants to be adsorbed. Weak interactions between the gas and the 3d/ 4s electrons keep the molecule in place while bonds are broken and formed. Some important heterogeneous catalysts are:

  • Fe used in the Haber process
  • Pt/ Rh used in catalytic converters. 
  • Ni used in hydrogenation. 

Transition METAL IONS act as homogeneous catalysts. They are able to change from one oxidation state to another during the reaction before returning to their original oxidation state. 

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Transition Metals- Formation of Complexes

A complex ion has a central metal ion which is surrounded by ligands. Ligands are molecules or anions with one or more lone pairs. An example of a complex is [Cu (H2O)6]2+.


[Cu (H2O)6]2+: The part in the bracket is the ligand

      The 6 is the coordination number, (the no. of dative covalent bonds)

      The 2+ is the overall charge on the ion. 

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Transition Metals- Formation of Complexes 2

Shapes of Complexes

Coordination number 6- octahedral

Coordination number 4- tetrahedral/ square planar

Coordination number 2- linear


1. How many ligands? (mono, di, tri, tetra, penta, hexa)

2. Name of ligands (alpha. order)- If the ligand is an ion add an '-o' eg. Cl-- chloro. 

H2O- aqua, NH3 -ammine~

3. Name metal- If negatively charged complex use Latin name, add an '-ate' eg. chromiumate.

~Iron- Ferrate, Copper- Cuprate~

4. Give oxidation state of central metal ion, eg. hexaaquachromium (III) ion

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Transition Metals- Formation of Complexes 3

Monodentate- Ligand molecule that bonds to a metal ion through a single bond

BidentateLigand molecule that bonds to a metal ion through two bonds.

PolydentateLigand molecule that bonds to a metal ion through more than one bond.

Chelate Rings- Occur when a bidentate/ polydentate forms a ring structure with a central metal ion. These rings make molecules more stable.

Ligand Substitution

Reactions when one ligand displaces another. They occur if the new complex formed is more stable than the previous complex.  

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Transition Metals- Formation of Complexes 4

Aqueous Copper (II) ions [Cu(H2O)6]2+                       blue solution

Aqueous Iron (III) ions [Fe(H2O)6]3+                        yellow solution

Aqueous Iron (II) ions [Fe(H2O)6]2+                        green solution

Copper (II) Ammonia Complex [Cu(NH3)4]2+ or [Cu(NH3)4(H2O)2]2+   deep blue solution

 Copper (II) chloride complex [CuCl4]2-                             yellow solution

Iron (II) hydroxide [Fe(OH)2]                            green precipitate

Iron (III) hydroxide [Fe(OH)3]                            rust brown/ orange precipitate

Copper (II) hydroxide  [Cu(OH)2]                            blue precipitate

The stability of a complex is obtained from its KSTAB value= complex stability constant. The bigger the number, the more stable the complex. 

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Transition Metals- Formation of Coloured Compounds

Transition metal ions are coloured because their 3d subshell electrons can be excited. When the transition metal is surrounded by ligands, the 5 d orbitals split into two different energy levels. The small energy gap between these orbitals allows visible light to be absorbed. When visible light falls on a coloured substance, the absorbed light is in the energy range that causes electronic transitions. The molecules become excited and move to higher energy levels before falling back to intermediate energy levels. Vibrational energy also increases. The colour we see is the complementary colour of that absorbed. 

Electrons in the lower level can be excited to a higher one. The energy required for this depends on the absorption of visible light. This can only happen in ions with a partially filled d subshell. Factors that affect this are:

  • the type of ligand
  • the shape of the complex
  • the coordination number of the complex
  • the charge on the central metal ion. 
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Transition Metals- Formation of Coloured Compounds


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Transition Metals- Formation of Coloured Compounds

Complementary Colours are opposite each other on the colour wheel.


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Electrode Potentials

Electrode Potentials allow you to decide which half equation should be oxidised and which should be reduced. Electrode Potential is a measure of the potential of a metal to release electrons and form ions. 

Electrochemical Cells

The potential difference between the 2 half cells is called Ecell value, and is measured by a voltmeter. When two half cells are joined:

  • The more negative electrode becomes the negative terminal of the cell and vice versa.
  • Electrons flow from negative terminal to the positive terminal. 

A salt bridge provides an ionic connection between two half- cells. It is usually made from a strip of filter paper soaked in potassium nitrate.

Standard Electrode Potentials

The standard hydrogen half cell is used as a reference electrode against which all other electrode potentials are measured. It is under standard conditions: 298K, 1atm, 1.00moldm-3  and is defined as 0.00V.

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Electrode Potentials 2

Standard Hydrogen Half Cell

H+(aq) + e- --> ½ H2(g)


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Electrode Potentials 3

The standard electrode potential Eɵ, of a half cell is defined as the potential difference between it and a standard hydrogen half cell. It is always positive. An electode potential chart has the most positive value at the bottom

Eɵcell values=  Eɵ[most postive electrode] - Eɵ[most negative electrode]

The most positive half cell/ least reactive is REDUCED

The most negative half cell/ most reactive is OXIDISED


  • We can predict feasibility of any redox reaction
  • The possibility of a reaction NOT the rate of reaction
  • Reactions only under standard conditions.
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