The d block: Characteristics of transition metals

The d block consists of three periods = 4,5 and 6 each contain ten elements

Across the first row of the d block, each element has one more proton in the nucleus and one more electron than the previous element. Each additional electron enters the 3d shell, not the outermost shell because the 4s orbital has already been filled.

The core is the electronic configuration of Argon which is the same for all atoms of the first row.This is 1s²2s²2p63s²3p6

These elements therefore essentially have the same outer electronic configuration as each other. They do not differ by a complete electron shell but by having one more electron in the inner, incomplete 3d sub-shell.

In the ground state of an atom, electrons are always arranged to give the lowest total energy.

The orbital energies in Chromium are such that putting one electron into each 3d and 4s orbital gives a lower energy than having two in the 4s orbital.

For Copper, putting two electrons in the 4s orbital would give a higher energy than filling the 3d orbitals.

Characteristic Properties: > Coloured Compounds > Variable oxidation states

The charcteristic properties of d-block elements are due to the presence of an inner incomplete d sub-shell. Electrons from both the inner d sub-shell and the outer s sub-shell can be involved in compound formation.

Not all d-block elements have an incomplete d sub-shell e.g. Zinc has 10 electrons in the d sub-shell so it is complete and Zinc and its 2+ ion do not show characteristic properties.

The only ion formed from Scandium (Sc3+) has no 3d electrons, so does not have charcteristic properties of transition metal ions.

A Transition Metal = An element which forms at least one ion with partially filled sub-shell of d electrons.

When d block elements form ions, it is the s electrons that are lost first, so ions of d-block elements contain only d electrons in their outer shell.

> All Metals > Similar to each other but distinctly different from s-block and p-block metals.

> Dense metals with high boiling and melting points

> Tend to be hard and durable > high tensile strength > good mechanical properties

Properties as a result of strong metallic bonding between the atoms in the metal lattice.

Often used for; cars, cooking utensils, building construction, bridges.

Transition metals can release electrons into the pool of mobile delocalised electrons from both the outer and inner shells.

More often used in the form of alloys - particularly with Iron > usually makes the metal harder and less malleable.

When a force is applied to a metal crystal, the layers of atoms can 'slide' over one another. Known as a slip. After slipping the atoms settle again into a close-packed structure. This makes metal malleable and ductile.

In alloys the different size atoms disrupt the orderly arrangement of atoms in the lattice and make it more difficult for the layers to slide over one another = harder and less malleable.

Smaller atoms, often non-metals can fit into the holes between the metal atoms - eg in Steel.

The non-metals such as Carbon or Nitrogen distort the lattice, making slip between layers more difficult.

In some steels, the Carbon forms Iron Carbide crystals which are very hard - added to the softer Iron regions this makes the steel very strong.

> Formation of compounds in a variety of Oxidation states > The catalytic activity of the elements and their compounds

> Strong tendency to form complexes > Formation of coloured compounds

Ionisation enthalpy = Energy needed to remove a aprticular electron from the atom or ion.

For Calcium, the first and second ionisation enthalpies are relatively low because this involves removing two s electrons from the outer shell. Then there is a sharp increase to the third ionisation enthalpy because it is much more difficult to remove further electrons from the 3p sub-shell.

For Vanadium there is a gradual increase in successive ionisation enthalpies because first the 4s electrons and then the 3d electrons are removed.

For all metals apart from Copper the sum of the first two ionisation enthalpies is low eneough for two electrons to be removed. (Copper can have the OS +1)

Except for Scandium and Titanium, all the elements show an OS +2, in which 4s electrons have been lost by ionisation.

All elements except Zinc show an OS of +3. The number of Oxidation States shown by an element increases from Sc to Mn. In these elements, the highest OS is equal to the total number of 3d and 4s electrons in atoms of the metal.

After Mn, the number of oxidation states begins to decrease and the increasing nuclear charge binds the d electrons more strongly, so they are harder to remove.

Lower OS tends to mean it is a simple ionic compound. Higher OS = bound covalently to an electronegative compound such as O or F.

Stabilitly of Oxidation States

> Higher OS less stable relative to lower OS (L>R)

>Compounds containing metals in high OS tend to be oxidising agents.

> The relative stability of the 2+ state with respect to the 3+ state increases across the series. For elements early in the series, elements with the 2+ state tend to be strong reducng agents. Later in the series the 3+ state is highly oxidising.

Heterogeneous - Solid metal catalyst & reactants in the gas or liquid phase

The electrons available in the 3d and 4s shells can form weak bonds to the reactants which break when the reaction has occurred and release the product.

Homogeneous - Usually in the aqueous phase with transition metals.

Transition metal ion forms an intermediate compound with one or more of the reactants - which then breaks down to form the products.

Transition metal ions = particularly effective catalysts in redox reactions - they can readily move from one OS to another