Chemistry · Ch 3 — Classification of Elements and Periodicity in Properties

The d-Block Elements (Transition Elements)

3.6.3

The d-Block Elements (Transition Elements)

The d-block elements occupy Groups 3 through 12 in the centre of the Periodic Table. Their defining feature is the filling of inner d orbitals with electrons — specifically, the $(n-1)d$ subshell. This is why they are called d-block elements.

The general outer electronic configuration for these elements is $(n-1)d^{1-10} ns^{0-2}$ . There is one notable exception: palladium (Pd) has the configuration $4d^{10} 5s^0$ , where the 5s orbital is empty. All d-block elements are metals.

Note

The term "transition element" is sometimes used interchangeably with "d-block element," but strictly speaking, transition elements are those that have partially filled d orbitals in their ground state or in any of their common oxidation states. This distinction matters for zinc, cadmium, and mercury, as we will see.

These metals share several characteristic properties:

They mostly form coloured ions.
They exhibit variable valence (multiple oxidation states).
They show paramagnetism (attraction to magnetic fields due to unpaired electrons).
They are often used as catalysts.

However, three elements at the end of each row — zinc (Zn), cadmium (Cd), and mercury (Hg) — have the electronic configuration $(n-1)d^{10} ns^2$ . Because their d subshell is completely filled, they do not show most of the typical transition element properties. For example, Zn²⁺ is colourless and diamagnetic (no unpaired electrons), and Zn shows only one common oxidation state (+2).

Watch out

A common mistake is to treat Zn, Cd, and Hg as full transition elements. They are d-block elements, but they are not transition elements in the strict sense because their d orbitals are fully filled in both the ground state and all common oxidation states.

The name "transition elements" comes from their position in the Periodic Table. They form a bridge — a transition — between the chemically active s-block metals on the left and the less active elements of Groups 13 and 14 on the right. This bridging role is reflected in their intermediate chemical behaviour.

General Electronic Configuration

The outermost shell (n) and the penultimate shell (n-1) are both involved in the electronic configuration. The general formula is:

(n-1)d^{1-10} ns^{0-2}

The superscripts indicate the range of electrons possible in each subshell. For example, scandium (Sc, atomic number 21) has the configuration $3d^1 4s^2$ , while zinc (Zn, atomic number 30) has $3d^{10} 4s^2$ .

\text{General configuration: } (n-1)d^{1-10} ns^{0-2}

The exception is palladium (Pd, atomic number 46), which has the configuration $4d^{10} 5s^0$ — the 5s orbital is empty, and the 4d subshell is completely filled.

Properties of d-Block Elements

The section lists several key properties. Each is explained below with its underlying reasoning.

1. They are all metals

All d-block elements are metals. This means they are good conductors of heat and electricity, are malleable and ductile, and have high melting and boiling points. The metallic character arises from the presence of delocalised electrons in the d orbitals, which can move freely through the metallic lattice.

2. They mostly form coloured ions

The colour of transition metal ions is a direct consequence of the partially filled d orbitals. When a transition metal ion is formed, the d orbitals split into two sets of different energies due to the electric field of surrounding ligands (this is called crystal field splitting). Electrons in the lower-energy d orbitals can absorb visible light and jump to the higher-energy d orbitals. The colour we see is the complementary colour of the light absorbed — that is, the colour of the light that is not absorbed.

Note

For example, the $\ce{[Cu(H2O)6]^{2+}}$ ion appears blue because it absorbs light in the red-orange region of the spectrum. The exact colour depends on the metal ion, its oxidation state, and the nature of the ligands attached.

Zinc, cadmium, and mercury ions (e.g., Zn²⁺, Cd²⁺, Hg²⁺) have a completely filled $d^{10}$ configuration. There are no empty d orbitals of the same energy to which an electron can jump, so no visible light is absorbed. Hence, these ions are colourless.

3. They exhibit variable valence (multiple oxidation states)

Transition metals can show a wide range of oxidation states because the energy difference between the $(n-1)d$ and $ns$ orbitals is small. This means that electrons from both the d and s subshells can participate in bonding. For example, manganese (Mn) shows oxidation states from +2 to +7.

Tip

A useful way to remember the possible oxidation states is to note that the maximum oxidation state often equals the total number of electrons in the $(n-1)d$ and $ns$ orbitals. For example, Mn has configuration $3d^5 4s^2$ , so it can lose up to 7 electrons (oxidation state +7).

In contrast, zinc, cadmium, and mercury have a filled $d^{10}$ subshell. The d electrons are held too tightly to be removed easily, so these elements typically show only the +2 oxidation state (by losing the two ns electrons). Mercury can also show the +1 oxidation state in the $\ce{Hg2^{2+}}$ ion, but this is an exception.

4. They show paramagnetism

Paramagnetism arises from the presence of unpaired electrons. When a substance is placed in a magnetic field, the unpaired electrons align with the field, causing the substance to be attracted to the field. The strength of paramagnetism is proportional to the number of unpaired electrons.

Transition metal ions with partially filled d orbitals often have unpaired electrons. For example, $\ce{Fe^{3+}}$ has five unpaired electrons (configuration $3d^5$ ), making it strongly paramagnetic.

Watch out

Do not confuse paramagnetism with diamagnetism. Diamagnetic substances have all electrons paired and are weakly repelled by a magnetic field. Zinc ions (Zn²⁺, $d^{10}$ ) are diamagnetic because all d electrons are paired.

5. They are often used as catalysts …

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