15 THE TRANSITION METALSWhat is the difference between a d-block element and a transition element?

Clue: Sc and Zn are not transition elements (see next slide)

Write the electronic configurations of the transition elements in period 4..

Sc 1s2 2s2 2p6 3s2 3p6 3d1 4s2

Ti [Ar] 3d2 4s2

Zn [Ar] 3d10 4s2

Why are the configurations of Cu and Cr as they are?

These elements are all d-block elements because the last electron that they received went into a d-orbital.

However, the definition of a transition element depends on its ability to form one or more stable ions with a partially filled d-orbital.

This property has consequences that affect the chemistry of these elements. Write the electronic configurations of the stable ions of scandium and zinc…

Look at p108 and list the characteristics of transition metals and their compounds…

Are they reactive?

Atomic radii, ionisation energies and reactivity

The d-orbital is inside the fourth shell, and so the radii of the atoms vary slightly across the 1st series

Although there is a slight decrease in radius due to increased nuclear charge, the screening from the inner shells reduces the effects of this (why are there little increases for Mn and Zn?)

The first and second ionisation energies vary little, and the differences are not significant. Compare successive IEs of Ca and Fe…

The “effective nuclear charge” increases, causing an increase in IE. Note the large jump for Ca.

With Fe, the extra ionisation energy is compensated more or less by the extra lattice enthalpy or hydration enthalpy evolved when the 3+ compound is made. The enthalpy of formations of FeCl2 or FeCl3 aren’t really that different.

The hydration of the ions formed by transition metals and the formation of complex ions and oxoions leads to the variable valency of transition metals.

This is because these processes are exothermic and compensate the successive ionisation energies.

The chemistry of the elements is less varied than the p-block elements in the same period due to the filling of the inner d-orbitals.

The d-orbitals get more involved in bonding across each of the transitions series above but by the end of each series the increases in nuclear charge means that these electrons are held more tightly and are less available for bonding

Text book pp108-109

COMPLEX IONS

These are central metal ion surrounded by neutral molecules or anions called “ligands”.

Ligands always have a lone pair of electrons that they can donate to the central ion and form dative covalent bonds.

The total number of dative bonds around the central metal ion is called the COORDINATION NUMBER.

Often, 6 bonds are formed in an octahedral arrangement.

Monodentate Ligands

H2O “aqua” CN- “cyano” NH3 “ammine” OH- “hydroxo”

Ions

Fe2+/Fe3+ “ferrate” Cu2+ “cuprate”

The overall charge on the ion is the sum of the charges of the central ion and the ligands, and the name given to the complex ion includes the type of ligand and the oxidation state of the central transition metal ion.

Eg.

The Hexaamminecobalt(III) ion

Name the above

Other possibilities…

Name the above

A bidentate ligand…

Chelation: the formation of multiple dative bonds between a polydentate ligand and the central metal atom/ion

Hydrazine cannot do this. Why?

EDTA - ethylenediaminetetraacetic acid: A polydentate ligand

This is a great example of a multi-dentate ligand…

Carbon monoxide acts as a better ligand to the central Fe ion and so Ligand Exchange occurs.

Carbon monoxide poisoning

Text book pp110-111/114-121

Transition metal ions and colour

Remember that the 3 d-orbitals are not identical:

When the ligands form dative covalent bonds with the central ion, the energy of all of the orbitals increases due to the greater electron density in general.

However, the orbitals that coincide with the bonds have their energy increased to a greater extent.

This can be represented as energy levels…

Transition metal ions must have a partially filled d-orbital, so that the absorption of light causes an electron to jump up to the higher energy level (excited state), and then return to the ground state releasing electromagnetic radiation which happens to be in the visible range for transition metals.

This explains why Zn and Sc are not transition elements, and do not form coloured compounds (although they do form complexes).

Cu+ also forms colourless complexes but Cu2+ no (it forms greens and blues).

When white light passes through a solution of a transition metal ion, some wavelengths of light are absorbed and the rest are reflected.

The colour wheel helps us to determine which colour will be seen

A colour change is observed when:• Oxidation number changes• Type of ligands change• The coordination number changes

Text book pp112-113

Types of reaction that transition metals undergo…

• Redox reactions• Acid-Base • Ligand exchange• Coordination number change

These reactions often cause a change in colour of the metal ion complexes

Hydration, deprotonation (acid-base) and ligand exchange

The hydrated ions can now be conceived of as the complex ion…

[Fe(H2O)6]3+

Aqua ions of transition metals are often acidic…

[Fe(H2O)6]3+ + H2O [Fe(H2O)5OH]2+ + H3O+

This is an acid-base or deprotonation reaction. How could we shift this equilibrium to the right?

Eqns:

The intense and distinct colours of the hydroxides formed can be used to identify the metal ion in question. The same precipitates can be observed when we add aqueous ammonia solution, due to the presence of OH-

Some transition metal ions in solution also show further reactions with aqueous alkali and ammonia solutions.

If we add a little NaOH(aq) or NH3(aq) to a green(violet) solution of hexaquachromium(III) ions, a green precipitate of Cr(H2O)3(OH)3 (s) forms.

In this case, the metal hydroxide formed must be amphoteric as it will react with XS NaOH(aq) and with H+(aq) ions…

Cr(H2O)3(OH)3 (s) + 3H+ (aq)

Cr(H2O)3(OH)3 (s) + 3OH- (aq)

We can shift these equilibria in either direction by adding acid or alkali

Ammonia is not a strong enough base to cause the metal hydroxides to show their amphoteric properties, but with Co2+, Ni2+ and Cu2+, a soluble ammine complex is formed…

Eg: Cu2+

This final reaction is an example of ligand exchange. Think about what complex ions formed would look like.

Water molecules are replaced when a more competitive ligand is added to the solution, eg. Aqueous Copper (II) ions and Conc HCl

Similar reactions occur with Cobalt (II) ions

Ligand exchange can be explained in terms of stability. The formation of dative covalent bonds is favoured thermodynamically, and the ligands that can donate their lone pair more efficiently will form complexes of lower energy.

Text book pp122-123

Other possible factors: • Charge on the ligands • Size of the ligands • Entropy • Orbitals available for bonding

Complexes are also formed by Be, Mg and Al due to their high charge density, although they are not coloured. Why?

Cu2+ Fe2+ Fe3+ Co2+ Cr3+

Formula of Aqueous ion

On adding a small amount of NaOH(aq)

On adding XS of NaOH(aq)

On adding a small amount of NH3(aq)

On adding XS of NH3(aq)

On adding conc HCl

Give the formulae and the colours for the complex ions formed in the following reactions…

LEARN!

Chromium

As well as the ligand exchange and acid-base reactions that we have already seen with Chromium, this element also displays a range of oxidation numbers and so REDOX reactions abound…

What is the formula of the dichromate(VI) ion?

Use your table of SEPs to show what happens if we add powdered Zn to a test tube containing this aqueous ion in acidic solution?

Give ionic equations and colours

We can use alkaline Hydrogen Peroxide solution to oxidise Chromium(III) Hydroxide to the chromate(VI) ion.

Oxidation of Cr(III)

Write the ionic equations and show that the reaction is feasible using SEPs

What happens if we acidify the resulting solution?

Is this redox?Text book pp128-129

Oxidation State

Ion  Colour Redox Equation Eo Values Compounds

+5 VO3-

VO2+

YELLOW VO2+(aq)  + 2H+(aq) + e-

VO2+(aq)  + H2O(l)+1.00 V

+4 VO2+ BLUE VO2+(aq + 2H+(aq) +e-

V3+(aq) + H2O(l)  +0.34 V

+3 V3+ GREEN V3+(aq) + e- V2+

(aq)     - 0.26 V

+2 V2+ VIOLET    

Vanadium is a great example of variable valency and coloured ions…

The +5 oxidation state is oxidising, and the +2 is reducing (it will even reduce water!!).

The +4 state is the most stable in the presence of air, and the +3 in its absence.

Vanadium

Write the half equations and the overall equation for the mild reduction of V(+5) to V(+4) using sulphite ions in acidic solution.

Write the half equations and the overall equation for the reduction of V (+5) by Zn…

Text book pp130-131

TRANSITION METALS AS CATALYSTS

They are affected by the reactions they catalyse, they do change form but they are not used up. An alternative route of lower energy is provided for the reaction.

There are two types: HOMOGENEOUS (same state as reactants) and HETEROGENEOUS (solids, usually).

Transition metals use their varied oxidation states to catalyse other reactions.

Eg. 2Ce4+ (aq) + Tl+ (aq) 2Ce3+ (aq) + Tl3+ (aq)

The reaction is slow, but can be catalysed with aqueous Mn2+ ions. How?

Heterogeneous catalysts work by absorbing the reactants onto their surface (active sites), thus weakening the bonds present and favouring the reaction to form the product which is then released (desorbed) from the surface.

This process depends on:

• the size of the lattice of the catalyst

• the strength of the bonds formed with the reactants and the products (must be strong and weak respectively)

• the use of available d-orbitals to form these bonds.

Catalysts are very specific for given reactions, for the above reasons.

Catalytic Converters Is this a heterogeneous or homogeneous catalyst in this case?

What are the reactions taking place? Describe the process…

What are the usual metals used to make these catalysts?

Look at P132 of the textbook and outline how Vanadium(V)Oxide acts as a catalyst in the contact process

Text book pp132-133

Homogeneous catalysis

Catalysts that are in the same phase as the reactants

The reaction occurs at a very slow rate, due to the fact that the 2 ions repel each other. Both Fe(II) ions and Fe(III) ions can act as catalysts for the reaction. Try and write the equations for these reactions showing the intermediate species

2 MnO4-(aq ) + 5 H2C2O4(aq ) + 6 H3O+(aq ) --> 2Mn2+(aq ) + 10 CO2(aq ) + 14 H2O

As we have seen before, Manganese in its various oxidation states can act as an autocatalyst in redox reactions. Remember the reaction with ethanedioic acid?

Considering that the 2 ions involved in this redox reaction are both negative, try to derive equations to show how the Mn2+ ions that are initially formed (slowly) speed up the reaction by being oxidised to Mn3+

Text book pp134-135