IB Chemistry - Periodicity

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Syllabus ref: 13.1

The first row of the 'd' block elements excluding scandium and zinc are known as transition metals. They share common characteristics that arise from having similar atomic and ionic radii.

Definition of transition metals

The term transition originally meant the elements that formed the transition between the 's' block and the 'p' block of the periodic table. However, it has since come to mean the elements which have compounds containing partiallly filled 'd' orbitals.

By this definition, in the first row of the transition metals the elements scandium and zinc are not included:

Scandium, the element, has an electronic configuration [Ar] 4s2 3d1, but it only forms ions with a 3+ charge, with a configuration of [Ar] 4s0 3d0.

Zinc, on the other hand, has a configuration with a full set of 'd' orbitals [Ar] 4s2 3d10, and only forms 2+ ions, in which the electronic arrangement is [Ar]4s0 3d10. Hence, scandium and zinc do not display many of the properties characteristic of the transition metals and are not included.

[Ar] 4s2 3d1
[Ar]4s2 3d10


Physical properties

The elements are hard metals with high melting points, a characteristic of strong metallic bonding. The transition metals have similar physical properties.

The '3d' sub-shell is inside the 4s sub-shell, meaning that as it gets filled moving from element to element, the inter-electron repulsion shields the outer 4s electrons from the increased nuclear charge. The consequence is that the atomic radius does not change as much as when crossing a short period, such as period 2 or 3.

Atomic radius initially shows a decrease from titanium to chromium, but once the electrons start to pack into the inner '3d' sub-shell the radius remains fairly constant.


Electronic configurations

The 4s orbitals fill up before the 3d orbitals, according to Hund's rule and the Aufbau principle, but there are two exceptions in the cases of chromium and copper.

Chromium half-fills its 'd' orbitals by promoting an electron from the 4s orbital. This gives rise to a half-filled 'd' shell that has some inherent stability, or at least enough extra stability to compensate for the energy needed to promote the 4s electron.

Copper similarly promotes one 4s electron to achieve a full set of 'd' orbitals.

[Ar]4s2 3d2
[Ar]4s2 3d3
[Ar]4s1 3d5
[Ar]4s2 3d5
[Ar]4s2 3d6
[Ar]4s2 3d7
[Ar]4s2 3d8
[Ar]4s1 3d10


Ion formation

Transition elements form positive ions by loss of electrons. The first electrons lost are the 4s electrons. Although this may seem counter intuitive as the last electrons to enter were the 3d, it seems as though the act of filling the '3d' sub-shell lowers its energy allowing the '4s' electrons to be lost first.

Example: Show the electronic configurations of an Fe atom and the Fe2+ ion.

  • Fe atom - [Ar] 4s2 3d6
  • Fe2+ ion - [Ar] 4s0 3d6

The common oxidation state of the transition elements is +2, reflecting the loss of the two electrons from the 4s orbital. Chromium's most stable oxidation state is not +2, however, reflecting the fact that there is only one 4s electron.

Once the 4s electrons have been lost, further loss of 3d electrons can lead to higher oxidation states.

Example: Show the electronic configurations of vanadium and the vanadium +2, +3, +4 and +5 oxidation states.

  • V atom - [Ar] 4s2 3d3
  • V(II) - [Ar] 4s0 3d3
  • V(III) - [Ar] 4s0 3d2
  • V(IV) - [Ar] 4s0 3d1
  • V(V) - [Ar] 4s0 3d0


Chemical properties

The effect of the inner 'd' orbitals is shown in the chemical properties. Their ionisation energies and atomic radii are similar, leading to similarities in reactivity. They all form double plus ions, but one of the characteristic features of the transition elements is their ability to form different oxidation states.

Their chemical properties will be dealt with in greater detail in the relevant section.