Transition Metals

Transition metals

A transition metal is a metal that can have at least one oxidation state with an electron in the d-orbital. According to this definition Scandium and Zinc are not transition metals because both Sc3+ and Zn2+ do not have electrons in their d orbitals.

General Properties

Size

The size of the transition metals is more or less the same. This is due to the fact that the last shell is actually the s orbital which is being shielded by the d electrons. The increase in the electron count in the d-orbital is counteracted with an increase in the positive charge in the nucleus, and therefore the effect of the nuclear charge on the outer electrons found in the s-orbital would be more or less constant. This would result in a similar atomic radius for all the transition metals in the same period.

  • Atomic radius remains more or less constant, increasing in size only slightly.
  • Ionisation energy decreases by a marginally small amount due to the constant atomic radius.

Metallic character

Transition metals are metals, and therefore they have metallic bonding, giving them similar properties as all metals:

  • High melting point and boiling point
  • High densities
  • High electrical and thermal conductivities
  • High tensile strength
  • Good mechanical properties

Oxidation state

All transition elements lose the 2 electrons found in the s orbitals to produce a 2+ ion. Apart from this, they can also lose electrons from the d-orbital, since these would have similar ionisation energies. Normally it is easier to lose unpaired electrons, although the most stable oxidations states are the +3 for the transition metals on the left and the +2 for the transition elements on the right. Manganese can have a maximum oxidation state of +7 since it is 4s2 3d5 and therefore it has 5 unpaired electrons in the d-orbitals.

Catalytic properties

Heterogeneous catalysts

A heterogeneous catalyst is a catalyst that is in a different state to the reactants, for example, a solid in a solution mixture. This normally offers an adsorption surface where the reaction can take place, just like Iron in the preparation of ammonia.

Homogenous catalysts

A homogeneous catalyst is one which is in the same state as the reaction mixture. In this case, transition metals can use their variable oxidations states. One such reaction is the reaction between iodide ions and thiosulfate.

Complexes

Complex compounds are produced between transition metals or ions by receiving electrons from a ligand via a dative bond. The ligand can be a neutral molecule or a negatively charged ion.

There are two types of ligands:

Monodentate: can only form 1 dative bond with the metal ion

Polydentate: can form more than one dative bond with the metal ion

Writing the formula

The formula is written in a specific order:

  1. Symbol of metal ion
  2. Formula of negatively charged ligands
  3. Formula of neutral ligands.
  4. If there is an overall charge the complex is written inside a square bracket.

Naming complexes

A complex should be named as follows:

  1. Ligands in alphabetical order
  2. Name of metal
    1. Normal name if complex is positive or neutral
    2. Latin name ending in –ate of the compound is negative
  3. Oxidation number of the metal

Shapes

Complexes can have a coordination number of 2, 4 or 6. A co-ordination number of 2 would give rise to a linear molecule, a coordination number of 4 can be either tetrahedral or square planar while co-ordination number of 6 gives rise to an octahedral complex.

Isomerism

Cis/trans isomerism

Chirality

Ligand groups and anions

Chromium(III) chlorides display the somewhat unusual property of existing in a number of distinct hydrates, forming a series of [CrCl3−n(H2O)n]z+. The main hexahydrate can be more precisely described as [CrCl2(H2O)4]Cl.2H2O. It consists of the cation [CrCl2(H2O)4]+and additional molecules of water and a chloride anion in the lattice. Two other hydrates are known, pale green [CrCl(H2O)5]Cl2.H2O and violet [Cr(H2O)6]Cl3.

Coloured compounds

When transition metals form ligands the overlapping of the orbitals do not remain degenerate, producing 2 different energy states. This is due to the interaction between the orbitals of the ligands with the d-orbital of the transition metal.

Electrons can be excited from the lower energy state to the higher energy state, and when the electron de-excites itself a wavelength is given off which corresponds to the wavelengths of visible light. The colour given off depends on the energy difference between the two states, which in turn depends on:

  • The central metal or ion
  • The ligands.

Chromium

General Properties

Electronic configuration: [Ar] 4s1 3d5

Oxidation states: +2, +3, +6

Amphoteric properties

Chromium can react with both acids and bases

Cr + 2HCl –> CrCl2 + H2

Cr + OH –> 2[Cr(OH)4] + 3H2

The reaction of Chromium with sulphuric acid is a redox reaction in which the sulphuric acid produces SO2 while the chromium is oxidised to the +3 oxidation state.

2Xr + H2SO4 –> Cr2(SO4)3 + 6H2O + SO2

On the other hand, it will not react with nitric acid since this will passivate the metal, making it unreactive.

Cr3+

Chromium burns in air to produce the oxide.

2Cr + 3O2 –> 2Cr2O3 (green solid)

This oxide dissolves in hydrochloric acid to give the hexaquachromium (III) with the chloride ions as the balancing anions (giving rise to the isomerism as discussed previously)

These compounds can be differentiated by the amount of AgCl that is precipitated on the addition of AgNO3. The Cl in the complex will not be precipitated.

The hexaaquachromium (III) is acidic since the Chromium will pull the electrons towards it making one of the waters to lose an H+

Cr6+

Chromium (III) can be oxidised to Chromium (VI) using a strong oxidising agent like the peroxide. Chromate is mostly known for their ability to act as oxidising agents, in which they are reduced to chromium (III).

Chromium (VI) can exist in two forms, CrO42- of Cr2O72- with the former being found in alkaline conditions and the latter being found in acidic conditions.

Oxide

The oxide is prepared by the action of concentrated sulfuric acid on the salt.

K2Cr2O7 + 2H2SO4 –> 2KHSO4 + 2CrO7 + H2O

This oxide can then react with hydrochloric acid to create CrO2Cl2 which is an oxidising agent for the reaction of methylbenzene to benzaldehyde.

CrO3 + 2HCl –> CrO2Cl2 + H2O

Manganese

General Properties

Electronic configuration: [Ar] 4s2 3d5

Oxidation states: +2, +3, +6

Reaction with acids

The metal reacts with acids to for the Mn2+ ion. The ion produced then forms very stable hexaaqua complexes, which are slightly pink in colour. This is a very stable complex.

Mn + HCl –> MnCl2 + H2

Mn2+ + 6H2O –> [Mn(H2O)6]2+

Reaction with halogens

Reacting Manganese with halogens gives the Mn2+ ion.

Mn + Cl2–> MnCl2

Oxide

The oxide of manganese produced in combustion with air depends on the amount of Oxygen produced.

The reaction pathway goes like this:

Mn + O2 –> MnO              black solid

4MnO + O2 –> 2Mn2O3     brown solid

2Mn2O3 + O2 –> 4MnO2   brown solid

With the Mn (IV) being the most stable)

Mn4+

MnO2 can act as both an oxidising agent and a reducing agent,m since it can be oxidised to Manganese (VII) while it can also be reduced to Mn2+.

Oxidising agent:

MnO2 –> Mn2+

Cl –> Cl2

 

Reducing agent:

MnO2 –> MnO42-

ClO3–>Cl

 

MnO2 is also oxidised by heating in alkali in contact with air.

Mn6+

The manganite (VI) ion disproportionates to give Manganese (VII) and Manganese (IV)

MnO42- –> MnO2

MnO42- –> MnO4 _

 

Mn7+

Manganese (VII) is a very good oxidising agent being itself reduced to Manganese (II)

Iron

General Properties

Electronic configuration: [Ar] 4s2 3d6

Oxidation states: +2, +3

Rust

Rust is the reaction of iron, water and oxygen. All 3 are needed in order for rust to take place.

Mixed oxide

When iron is reacted with oxygen on heating or steam the mixed oxide is formed.

3Fe + 2O2 –> Fe3O4

3Fe + 4H2O –> Fe3O4 +4H2

Reaction with acid

Iron reacts with acids to form Iron(II) compounds since these are not very strong oxidising agents.

Fe + HCl –> FeCl2 + H2

Fe + H2SO4 –> FeSO4 + H2

Both of these salts are hydrated. Preparation of the anhydrous chloride by heating is not possible, while preparation of the sulphate is. This is due to the fact that the chloride would actually react with the hydroxide.

FeCl2.6H2O –> FeOHCl + 5H2O + HCl

Fe2+

FeS

Prepared by direct synthesis

Fe + S –> FeS

This happens because S2- is a strong reducing agent and it will reduce Fe3+ to Fe2+.

Fe(OH)2

Prepared by the addition of OH to any Iron(II) salt.

Fe + OH –> Fe(OH)2 green

FeX2

All the halides can be produced by the reaction of the iron with the acid.

Fe + 2HCl –> FeCl2 + H2

The bromide and the iodide can also be produced direct combination.

Fe + Br2 –> FeBr2

The iron should be in excess to ensure that no FeBr3 is formed.

FeCO3

Iron is acidic and dissolution of iron carbonate in water would give rise to iron hydroxide and the hydrogen carbonate.

[Fe(H2O)6]2+ + 2CO32- –> [Fe(OH)2(H2O)] + 2HCO3

FeSO4

Iron(II) sulphate is prepared by the reaction of iron with sulfuric acid.

The hydrated crystals are green and turn white when the water of crystallisation is removed.

On strong heating iron (III) oxide is produced together with a mixture of SO2 and SO3

2FeSO4 –> Fe2O3 + SO2 + SO3

Brown ring test

This is a test for nitrates.

NO3+ FeSO4 + H2SO4 –> brown ring

NO3 + 4H+ 3e–> NO + 2H2O

Fe2+ –> Fe3+ + e

NO3 + 4H+ + 3Fe2+ –> NO + 2H2O +3Fe3+

NO can act as a ligand and it will displace one of the water ligands on the iron, and this would form the brown ring visible in the test.

[Fe(H2O)6]2+ +NO à Fe(H2O)5NO]2+ + H2O

Fe3+

Oxide

The oxide os prepared by the decomposition of Iron (II) sulphate or by the reaction of iron (III) with hydroxide ions.

2FeSO4 –> Fe2O3 + SO2 + SO3

Fe3+ + OH –> Fe(OH)3 à Fe2O3.xH2O

Halides

The chloride and bromide are prepared by direct combination with the halogens. These are both collected and purified by sublimation.

2Fe + 3Cl2 –> 2FeCl3

2Fe + 3Br2 –> 2FeBr3

It is noted that Bromine has to be in excess in order to ensure that FeBr3 is formed.

Thiocyanate

Thiocyanate and Fe3+ can form a complex with a deep blood red colouring. The thiocyanate will displace a water molecule to form the complex.

[Fe(H2O)6]3+ + SCN–> Fe(SCN)(H2O)5]2+  + H2O

Copper

General Properties

Electronic configuration: [Ar] 4s1 3d10

Oxidation states: +1, +2

Reactions of Copper

It reacts with strong oxidising agents such as HNO3 and H2SO4 (both concentrated) to give nitrogen dioxide and sulfur dioxide respectively.

Cu + 4HNO3 –> Cu(NO3)2 + 2H2O + NO2

Cu + 4H2SO4 –> CuSO4 + 2H2O + SO2

Cu2+

Copper (II) oxide can be prepared from the decomposition of copper carbonate, copper nitrate or copper hydroxide.

2Cu(NO3)2 –> 2CuO + 4NO2 + O2

This oxide is basic and therefore it can be used to produce salts by reacting with the respective acids.

CuO + H2SO4 –> CuSO4 + H2O

Cu+

Less stable than copper (I)

Copper (I) disproportionates to give copper and Copper (II)

2Cu+ –> Cu + Cu2+

Iodine and Iodide only form the copper (I) salts.

Cu + I2 –> CuI

2Cu2+ + 4I –> 2CuI + I2

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