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Transition Metals and their Properties

 

A transition metal or element is generally defined as a metal which has partially filled d orbitals in the neutral atom or in any of its usual positive oxidation states.

The first transition series extends from scandium (Sc) to copper (Cu).

Electronic Configuration of The First Transition Series

21Sc - [Ar] 3d14s2
22 Ti - [Ar] 3d24s2
23V - [Ar] 3d3s2
24Cr - [Ar] 3d54s1
25Mn - [Ar] 3d54s2
26Fe - [Ar] 3d64s2
27Co - [Ar] 3d74s2 
28Ni - [Ar] 3d84s2 
29Cu - [Ar] 3d104s1 

Note:

*[Ar] represents the electronic configuration of argon (1s22s22p63s23p6), which comes before the outer electronic levels of transition elements.

Only the outer electronic configuration are shown above. However, always write the complete configurations whenever you are asked to write the electronic configurations of transition metals, example, Sc 1s22s22p63s23p63d14s2.

*All the elements have their d orbitals partially filled. Example, 27Co - 3d74s2 and 28Ni3d84s2 as shown below.

27Co3d74s2  

                                d-orbitals        
 

28Ni3d84s2  

Zinc has completely filled d orbitals, it is therefore not regarded as a transition metal.

*The electronic configuration of chromium is predominantly [Ar] 3d54s1 instead of [Ar] 3d44s2 as you might have expected.

The reason why this is the case is that, having five electrons in the d orbitals produces a kind of stability known as partial stability on the metal, considering that five is half of ten (for full stability).

Therefore, after vanadium, the next electron, together with one electron of 4s preferably go into d orbitals (to make it five). The same reason explains the configuration of copper as shown above.

Properties (Physical and Chemical) of Transition Elements

Transition metals show unique physical and chemical properties. These include:

Physical Properties

1. They are all Metals.

2. They are hard, malleable, ductile and are of very high melting and boiling points. Compared with the main group metals (such as metals of group one), transition metals have higher boiling and melting point .

The reason for this is the presence of very strong metallic bonding - due to large number of valence electrons involved in it.

3. They are good conductors of electricity. This is because their electrons are very free to move about within the available vacant d orbitals.

4. They exhibit paramagnetism. This is due to the presence of unpaired electrons.

Paramagnetism is the phenomenon whereby substances are weakly attracted into a magnetic field. Substances which are paramagnetic contain one or more electrons that are not paired off in the usual way.

Consequently, there is a magnetic moment which results in attraction by a magnetic field. The higher the number of unpaired electrons, the greater the paramagnetic properties.

Example, 27Co has three unpaired electrons as shown above, and is therefore more paramagnetic than 28Ni, which has two.

Note that a diamagnetic substance is repelled by a magnetic field. All its electrons are paired.

5.They may function in their metallic state or in certain of their compounds as catalyst.

This is possible because they possess free and readily available d orbitals (energy levels), which they provide for reducing the activation energy of the reactions they catalyze.

6. They are capable of forming alloys with one another, and with other elements.

Chemical Properties

1. Variable oxidation state

Apart from scandium which forms only +3 oxidation state, transition metals show various positive oxidation states by losing electrons from both their 4s and 3d orbitals (the energy difference between these orbitals is small).

Oxidation state of +2 is formed by the loss of the two electrons of the 4s (except in Cr and Cu, where the two electrons are lost from 4s and 3d orbitals, with each orbital losing an electron.

Compounds of the metals (from Ti to Cr) in the +2 state are ionic and are strong reducing agents (they are strongly oxidized). Example, Ti2+, V2+ and Cr2+.

On the other hand, the +2 state of the elements from Mn to Cu is more stable and less readily oxidized (they are less reducing agents.)

Elements from Ti to Mn show higher oxidation states by loss of higher number of electrons from both s and d orbitals. Example, Ti 4+ (as in TiCl4); Cr6+ (as in CrO42-) and Mn7+ (as in MnO-4).

Note:

*Ions in high oxidation states (such as Cr6+, Mn7+ or V5+) do not occur as independent units, but form covalent bonds with other units such as oxygen.

Example, in aqueous solution, Cr6+ forms CrO42-, V5+ forms VO3+ or VO43- and Mn7+ forms MnO4-. This is because, the higher oxidation states are extremely polarizing.

*The oxides of the metals in high oxidation states are acidic and their salts are oxidizing agents (example, KMnO4).

*There are three particularly stable electronic arrangements in the first transition series above:

Those with empty 3d orbitals (example, Sc3+ and Mn7+ 1s22s22p63s23p6)

Those with half-filled 3d orbitals (example, Mn2+ and Fe3+ 1s22s22p63s23p63d5).

Those with completely filled 3d orbitals (example Cu+ 1s22s22p63s23p63d10)

For example, Fe3+ is more stable than Fe2+ because its 3d orbitals are half-filled. Hence, Fe2+ is easily oxidized to Fe3+.

Mn7+ is more stable than Mn3+ because it has empty d orbitals. And Cu+ is more stable than Cu2+ because it has completely filled d orbitals.

The table below shows the different possible oxidation states of the first transition series:

Transition Metals

Sc

Ti

V

Cr

Mn

Fe

Co

Ni

Cu

Oxidation State 
 


 

+3

+2 +3 +4

+2 +3 +4 +5

+2 +3 +6

+2 +3 +7

+2 +3

+2 +3

+2 +3

+1 +2

 

2. Complex ion Formation

Transition metal ions have small cationic size, large ionic charge and empty or partially filled 3d orbitals, hence they are capable of accepting electron pairs from either charged or neutral particles into their d orbitals to form dative or coordinates bonds, resulting in complex ion formation.

The electron donors are called ligands (examples, NH3, H2O, and CN-, and ethylene diamine NH2CH2CH2NH2). The number of ligands attached to a metal ion is regarded as its coordination number and it is commonly 4 to 6.

Complex ion formation are aided by the following factors:

(a). Oxidation number - complex ion formation increases with increase in oxidation number. That is, the higher the oxidation number of the element, the greater the stability of the complex ion formed.

Example, hexamine cobalt(III) ion, [Co(NH3)6]3+ is more stable than hexamine cobalt(II) ion, [Co(NH3)6]2+ as it does not lose ammonia, while the cobalt(II) loses ammonia and gives the reaction of cobalt(II) in solution.

(b). Atomic number - the stability of a particular complex ion increases within a transition series with increase in atomic number. Example, the stability of complex ion formed between NH3 and a transition metal in the oxidation state of +2, from Mn to Cu is in the order:

Cu2+> Ni2+ > Co2+ > Fe2+

(c). The nature of ligands

The stability of complex ion formed also depends on the degree of interaction between the ligand and the central ion.

Some ligands form more stable complexes than others. Example, the following is the order of stability of certain ligands with the same transition metal ion (in the same oxidation state):

Cu(CN)42- > Cu(NH3)42+ > Cu(H2O)42+

3. Formation of coloured ions - transition metal ions tend to be highly coloured. This is due to the easy movement of electrons between the d orbitals, resulting in absorption or emission of light whose frequencies lie in the visible region.

The colour is usually associated with a specific oxidation state of a given element. The table below shows the colours of some complex ions:

Oxidation State

Example of Compound

Colour

Cu(I) 

CuCl

Colourless

Cu(II) 

CuSO4.5H2O

Blue

Mn(II) 

MnCl2, Mn(NO3)2

Colourless

Mn(III) 

Mn2O3

Violet

Mn(VII) 

KMnO4

Purple

Cr(II) 

CrCl2.6H2O

Blue

Cr(III) 

Cr2(SO4)3.18H2O

Green

Cr(VI) 

K2Cr2O7

Orange

Fe(II) 

FeSO4.7H2O

Green

Fe(III) 

FeCl3.6H2O

Yellow

FeCl3

Very dark or black

Co(II) 

CoCl2.6H2O

Pink

Ni(II) 

Ni(NO3)2.6H2O

Green

V(II) 

VSO4.6H2O

Violet

V(III) 

V(H2O)63+

Green

V(IV) 

V(H2O)64+

Blue

Ti(III) 

Ti(H2O)63+

Violet

Ti(IV) 

TiCl4.2H2O

Colourless

 

  

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