General Properties of Transition Metals

According to the Aufbau Principle, electrons occupy the lowest energy subshells first. In the first transition series, the $4s$ subshell is filled before the $3d$ subshell because it is at a lower energy level in neutral atoms.

There are two notable exceptions: Chromium ($[Ar] 3d^5 4s^1$) and Copper ($[Ar] 3d^{10} 4s^1$). In these cases, an electron is promoted from the $4s$ to the $3d$ subshell to create a half-full or full d-subshell, which provides extra energetic stability.

When transition metals form ions, they always lose electrons from the 4s subshell first. Once the $3d$ orbitals are occupied, the repulsion between electrons pushes the $4s$ subshell to a slightly higher energy state than the $3d$, making the $4s$ electrons the most 'outer' and easiest to remove.

Unlike s-block metals (like Sodium, which only forms $+1$), transition metals can exist in multiple stable oxidation states. This occurs because the energy levels of the $4s$ and $3d$ subshells are very close together.

The energy required to remove successive electrons is relatively small, allowing the metal to lose different numbers of electrons depending on the chemical environment and the nature of the reacting species.

For example, Manganese can exhibit oxidation states ranging from $+2$ to $+7$. Higher oxidation states are typically found in complex ions or compounds where the metal is bonded to highly electronegative elements like Oxygen.

Transition metal ions have a high charge density and vacant d-orbitals, which allows them to attract and bond with ligands. A ligand is a molecule or ion that donates a lone pair of electrons to form a coordinate (dative) bond with the central metal ion.

The resulting structure is called a complex ion. The number of coordinate bonds formed is known as the coordination number, which determines the geometric shape of the complex (e.g., octahedral for 6 bonds, tetrahedral for 4 bonds).

Common ligands include neutral molecules like $H_2O$ and $NH_3$, or negative ions like $Cl^-$ and $CN^-$. The ability to form these complexes is a defining chemical characteristic of the transition series.

Most transition metal compounds are brightly coloured. This arises because the presence of ligands causes the five d-orbitals to split into two different energy levels; when an electron absorbs light energy, it can jump from a lower to a higher d-orbital (d-d transition).

The color we perceive is the complementary color to the wavelength of light absorbed. For instance, if a complex absorbs red light, it will appear blue-green (cyan) to the observer.

Transition metals are excellent catalysts because they can change oxidation states to provide alternative reaction pathways (homogeneous catalysis) or provide a surface for reactants to adsorb onto (heterogeneous catalysis).

It is vital to distinguish between a d-block element and a transition metal. While all transition metals are in the d-block, not all d-block elements are transition metals.

• The 4s Rule: Always remember to remove $4s$ electrons before $3d$ electrons when writing ion configurations. This is the most common area where students lose marks.

• Definition Precision: If asked to define a transition metal, you must mention 'incomplete d-subshell' and 'stable ion'. Simply saying 'd-block' is insufficient.

• Stability Patterns: Recognize that $d^0$, $d^5$, and $d^{10}$ configurations are particularly stable. This explains the exceptions in Cr and Cu and the common oxidation states of other metals.

• Color Logic: If a solution is colourless, it likely contains an ion with a completely full ($d^{10}$) or completely empty ($d^0$) d-subshell, as d-d transitions are impossible in these states.