What distinguishes a transition metal from a representative metal?

Understanding the periodic table is fundamental to grasping the difference between transition and representative metals. Periodic table position is dictated by the atomic structure of elements, particularly their electron configurations. Representative metals are found in groups 1A and 2A, which are the alkaline metals and alkaline earth metals respectively, and also in groups 3A to 8A, encompassing elements such as the noble gases, halogens, and other nonmetals and metalloids.

On the other hand, transition metals are located in the central block of the table, specifically in groups 3 to 12. This area is also known as the 'd-block', because the defining characteristic of these metals is the filling of their d orbitals with electrons as one moves across the period. Since the properties of elements are largely determined by their electron arrangements, their location on the periodic table can reveal much about their chemical behavior.

Electron configuration is the key to understanding an element's chemical properties. The electron configuration of an atom describes the distribution of electrons in its atomic orbitals. Representative metals have their outermost electrons in s or p orbitals. For example, sodium (Na) has a single s electron outside a closed shell, while aluminum (Al) carries three p electrons beyond its inner closed shells.

In contrast, transition metals have partially filled d orbitals. This can result in a larger number of electrons that may participate in chemical bonding, differentiating them from the representative metals whose valence electrons are more strictly limited to the s and p orbitals. This unique electron distribution in transition metals is responsible for many of their distinctive properties, such as the ability to form various compounds with different oxidation states.

Closely linked to the electron configuration of an element is its tendency to exist in varying oxidation states. An oxidation state is an indicator of the degree of oxidation or reduction of an element in a chemical compound. Representative metals typically exhibit a limited number of oxidation states because they lose a predictable number of s- or p-orbital electrons to achieve a stable electron configuration. For example, magnesium (Mg) consistently shows a +2 oxidation state in its compounds.

Transition metals are unique in that they frequently exhibit a wide range of positive oxidation states. Elements like iron (Fe) can exhibit +2 and +3 oxidation states, among others, because they can lose varying numbers of d-orbital electrons in addition to their s electrons. This multiplicity of oxidation states contributes to the formation of a rich variety of compounds, allowing for great versatility in their chemical reactions.

One of the most visually striking aspects of transition metals is their ability to form coloured complexes. When transition metals form complexes with other molecules or ions, they often display a wide array of colours. This phenomenon occurs because of the d-orbital electrons in transition metals, which can absorb specific wavelengths of light to become excited to higher energy levels before falling back and emitting light.

This does not typically occur in representative metals since they lack the partially filled d orbitals requisite for such electronic transitions. The absorption of light highlights the difference between the full or empty p orbitals in representative metals and the partially filled d orbitals of transition metals. For example, copper(II) sulfate forms a bright blue complex, while the compounds of calcium, a representative metal, are generally colourless. Understanding these concepts helps to illuminate the roles transition metals can play in everything from industrial catalysts to biological systems.