Transition elements: electronic configuration, occurrence, and general trends, oxidation states, and catalytic behaviour

1. Introduction to Transition Elements

Transition elements are defined as elements that have partially filled d orbitals in their ground state or in any of their stable oxidation states. They are located in the p-block of the periodic table, specifically in groups 3 to 12. These elements are also known as d-block elements because their last electron enters one of the five d orbitals.

The general electronic configuration of transition elements is (n-1)d^1-10 ns^1-2. Here, 'n' represents the principal quantum number of the outermost shell, and (n-1) represents the penultimate shell where the d orbitals are being filled.

2. Occurrence of Transition Elements

Transition elements are found widely distributed in the Earth's crust, often in the form of oxides, sulfides, carbonates, and silicates. They are not found in their free state due to their reactivity, except for noble metals like gold and platinum.

Some common transition metals and their ores include:

• Iron (Fe): Hematite (Fe₂O₃), Magnetite (Fe₃O₄), Siderite (FeCO₃)
• Copper (Cu): Copper glance (Cu₂S), Malachite (CuCO₃.Cu(OH)₂), Cuprite (Cu₂O)
• Zinc (Zn): Zinc blende (ZnS), Calamine (ZnCO₃)
• Manganese (Mn): Pyrolusite (MnO₂)
• Chromium (Cr): Chromite (FeCr₂O₄)
• Nickel (Ni): Garnierite (Ni₂SiO₄.nH₂O)

The abundance of these elements in the Earth's crust varies. Iron is the fourth most abundant element, while copper, zinc, and nickel are found in much smaller quantities.

3. Electronic Configuration

The electronic configuration of transition elements is characterized by the filling of (n-1)d orbitals. The filling follows the Aufbau principle, Hund's rule, and the Pauli exclusion principle. However, there are some exceptions due to the extra stability associated with completely filled or half-filled orbitals.

The first transition series (3d series) includes elements from Scandium (Sc) to Zinc (Zn).

Element	Symbol	Atomic Number	Electronic Configuration
Scandium	Sc	21	[Ar] 3d1 4s2
Titanium	Ti	22	[Ar] 3d2 4s2
Vanadium	V	23	[Ar] 3d3 4s2
Chromium	Cr	24	[Ar] 3d5 4s1 (Exception)
Manganese	Mn	25	[Ar] 3d5 4s2
Iron	Fe	26	[Ar] 3d6 4s2
Cobalt	Co	27	[Ar] 3d7 4s2
Nickel	Ni	28	[Ar] 3d8 4s2
Copper	Cu	29	[Ar] 3d10 4s1 (Exception)
Zinc	Zn	30	[Ar] 3d10 4s2

Exception Explanation: Chromium ([Ar] 3d⁵ 4s¹) and Copper ([Ar] 3d¹⁰ 4s¹) exhibit exceptions to the general filling order. This is because a half-filled d-subshell (3d⁵) and a completely filled d-subshell (3d¹⁰) provide extra stability. The promotion of one electron from the 4s orbital to the 3d orbital in these cases leads to a more stable configuration.

The second transition series (4d series) includes elements from Yttrium (Y) to Cadmium (Cd), with general electronic configuration (n-2)f^0-14 (n-1)d^1-10 ns^1-2. The third transition series (5d series) includes elements from Lanthanum (La) and Hafnium (Hf) to Mercury (Hg). The fourth transition series (6d series) includes Actinoids and elements from Rutherfordium (Rf) to Copernicium (Cn).

4. General Trends in Transition Elements

Transition elements exhibit several characteristic trends across a period and down a group, which are a consequence of their electronic configurations.

4.1. Atomic and Ionic Radii

Across a period (from left to right), the atomic and ionic radii of transition elements generally decrease. This is because the nuclear charge increases with the addition of protons, pulling the electrons closer to the nucleus. However, the increase in the number of electrons in the (n-1)d subshell shields the outer ns electrons from the full effect of the nuclear charge. The shielding effect of d electrons is not as effective as that of s or p electrons. The decrease in radius is not very sharp because the added electron goes into the inner (n-1)d orbitals, which are shielded by the outer ns electrons. In the middle of the series, the radii tend to remain almost constant.

Down a group, the atomic and ionic radii generally increase. This is due to the addition of a new electron shell with each successive period. For example, the radii of the 4d series elements are generally larger than those of the corresponding 3d series elements, and the 5d series elements are larger than the 4d series elements.

A notable exception is the lanthanide contraction, where the radii of the 5d series elements are very similar to those of the corresponding 4d series elements. This is because the 4f orbitals are filled in the lanthanide series, and f electrons have poor shielding effects, leading to a significant decrease in size.

4.2. Ionization Enthalpy

The ionization enthalpy of transition elements generally increases across a period. This is due to the increasing nuclear charge and the relatively poor shielding effect of the (n-1)d electrons. However, the increase is not as regular as in p-block elements due to the small energy differences between the (n-1)d and ns electrons and the relatively stable configurations like d⁵ and d¹⁰.

The first ionization enthalpies of transition metals are generally higher than those of alkali and alkaline earth metals but lower than those of p-block elements.

Down a group, ionization enthalpies generally decrease due to the increase in atomic size and the shielding effect of inner electrons. However, the decrease is less pronounced than in the main group elements.

4.3. Oxidation States

Transition elements exhibit variable oxidation states. This is because they have electrons in both the outermost ns orbital and the inner (n-1)d orbitals. These electrons can be lost or shared during chemical bonding, leading to multiple oxidation states. The energy difference between the (n-1)d and ns electrons is small, making it possible for electrons from both subshells to participate in bonding.

The variable oxidation states are a hallmark of transition elements. For example, Manganese (Mn) exhibits oxidation states from +2 to +7.

The lowest oxidation state usually corresponds to the loss of ns electrons, while higher oxidation states involve the loss of (n-1)d electrons as well. The most common oxidation state for many transition metals is +2, corresponding to the loss of the two 4s electrons.

The stability of oxidation states varies. Oxidation states that lead to a half-filled (d⁵) or completely filled (d¹⁰) d-subshell are often more stable.

4.4. Magnetic Properties

Transition metal ions often exhibit magnetic properties due to the presence of unpaired electrons in their d orbitals. These unpaired electrons give rise to two types of magnetism:

• Paramagnetism: Occurs when an atom or ion has unpaired electrons. These substances are attracted to an external magnetic field. The magnetic moment ($\mu$) can be calculated using the spin-only formula: $\mu = \sqrt{n(n+2)}$ B.M., where 'n' is the number of unpaired electrons and B.M. stands for Bohr Magneton.
• Diamagnetism: Occurs when all electrons in an atom or ion are paired. These substances are weakly repelled by an external magnetic field.

Most transition metals and their compounds are paramagnetic because they have unpaired d electrons. However, if all d electrons are paired (e.g., in Zn²⁺, Sc³⁺, Ti⁴⁺, Cu⁺), the species will be diamagnetic.

4.5. Formation of Coloured Compounds

Many transition metal compounds are coloured. This colour arises from the absorption of certain wavelengths of visible light, which causes the excitation of an electron from a lower energy d orbital to a higher energy d orbital. This phenomenon is known as d-d transition. The colour observed is complementary to the colour of the light absorbed.

The presence of unpaired electrons is essential for d-d transitions. If a transition metal ion has no unpaired electrons (e.g., Zn²⁺, Sc³⁺), its compounds are usually colourless.

Examples:

• Cu²⁺ (aq) is blue because it absorbs orange light.
• Ni²⁺ (aq) is green because it absorbs violet light.
• Fe³⁺ (aq) is yellowish-brown.
• Mn²⁺ (aq) is pale pink.

The intensity and shade of the colour can depend on the oxidation state of the metal, the nature of the ligands, and the geometry of the complex.

4.6. Formation of Complex Compounds

Transition metals have a strong tendency to form complex compounds. This is due to several factors:

• Small size of metal ions: Transition metal ions are generally small, leading to a high charge density.
• High charge on metal ions: Many transition metals can exist in higher oxidation states, resulting in a high charge.
• Presence of vacant d orbitals: The presence of vacant (n-1)d orbitals allows them to accept lone pairs of electrons from ligands, forming coordinate bonds.

Ligands are molecules or ions that donate a lone pair of electrons to the central metal ion to form a coordinate bond. Examples of complex compounds include [Cu(NH₃)₄]²⁺, [Fe(CN)₆]^4-, and [Ni(dmg)₂].

4.7. Formation of Alloys

Transition metals readily form alloys with other metals. An alloy is a mixture of two or more elements, at least one of which is a metal, and the resulting material has metallic properties. Transition metals have similar atomic radii and crystal structures, which allows them to form solid solutions where one metal atom can substitute for another in the crystal lattice.

Examples of alloys involving transition metals:

• Steel: An alloy of iron and carbon.
• Bronze: An alloy of copper and tin.
• Brass: An alloy of copper and zinc.
• Stainless Steel: An alloy of iron, chromium, and nickel.

4.8. Metallic Properties

Transition elements are generally hard, strong, and possess high melting and boiling points. This is due to the strong metallic bonding arising from the delocalization of valence electrons. The involvement of (n-1)d electrons in metallic bonding contributes to their strength and high melting points.

They are good conductors of heat and electricity. Most transition metals are lustrous and can be hammered or drawn into wires (malleable and ductile).

5. Oxidation States of Transition Elements

The ability of transition elements to exhibit multiple oxidation states is one of their most characteristic properties. This arises from the participation of both ns and (n-1)d electrons in chemical bonding.

General Trends in Oxidation States:

• The lowest oxidation state is usually +1 or +2, corresponding to the loss of ns electrons.
• As we move across the series, the number of available electrons for bonding increases, leading to an increase in the number of possible oxidation states.
• The highest oxidation state usually corresponds to the loss of all ns and (n-1)d electrons. For example, in Manganese (3d⁵ 4s²), the highest oxidation state is +7.
• The stability of oxidation states varies. For the first transition series, the +2 oxidation state is common. However, as the number of d electrons increases, higher oxidation states become more prevalent.
• In the middle of the series (e.g., Mn, Fe), oxidation states tend to be more stable due to the involvement of a larger number of electrons.
• At the end of the series (e.g., Cu, Zn), the oxidation states tend to decrease. Zinc almost exclusively shows +2 oxidation state because its electronic configuration is [Ar] 3d¹⁰ 4s², and losing the 4s electrons results in a stable, completely filled 3d subshell.

Example: Manganese (Mn)

Electronic configuration: [Ar] 3d⁵ 4s²

Oxidation states: +2, +3, +4, +5, +6, +7

• Mn²⁺: [Ar] 3d⁵ (half-filled, stable)
• Mn³⁺: [Ar] 3d⁴
• Mn⁴⁺: [Ar] 3d³
• Mn⁶⁺ (in MnO₄^2-): [Ar] 3d¹
• Mn⁷⁺ (in MnO₄^-): [Ar] 3d⁰ (highest, stable in oxidizing conditions)

Relating Oxidation States to Electronic Configuration:

A useful mnemonic is that the oxidation state of a transition element is often equal to the sum of the number of electrons in the ns and (n-1)d orbitals, up to a certain point.

6. Catalytic Behaviour of Transition Elements

Transition metals and their compounds exhibit excellent catalytic activity. This is a crucial property that makes them indispensable in many industrial processes. The catalytic activity arises from several factors:

• Variable Oxidation States: Transition metals can exist in multiple oxidation states. This allows them to undergo oxidation and reduction reactions, facilitating the interaction with reactants. For example, in a catalytic cycle, a metal can adsorb a reactant, change its oxidation state, and then release the product.
• Formation of Intermediate Complexes: Transition metals can form intermediate complexes with the reactants. These intermediates have lower activation energy, thus speeding up the reaction.
• Surface Area: Many transition metals are used in finely divided form or as supported catalysts, which provides a large surface area for the reaction to occur. This increases the rate of reaction.
• Unpaired Electrons: The presence of unpaired electrons on the d orbitals can also play a role in activating reactants or stabilizing transition states.

Examples of Catalytic Activity:

• Contact Process for Sulfuric Acid Production: Vanadium(V) oxide (V₂O₅) is used as a catalyst to oxidize sulfur dioxide (SO₂) to sulfur trioxide (SO₃):

  2SO₂(g) + O₂(g) $\xrightarrow{V_2O_5}$ 2SO₃(g)

• Haber Process for Ammonia Synthesis: Iron (Fe) is used as a catalyst for the synthesis of ammonia from nitrogen and hydrogen:

  N₂(g) + 3H₂(g) $\rightleftharpoons$ 2NH₃(g)

• Hydrogenation of Unsaturated Hydrocarbons: Nickel (Ni), Palladium (Pd), and Platinum (Pt) are used as catalysts to hydrogenate alkenes and alkynes:

  C₂H₄(g) + H₂(g) $\xrightarrow{Ni}$ C₂H₆(g)

• Oxidation of Ethylene to Ethylene Oxide: Silver (Ag) acts as a catalyst.
• Ziegler-Natta Catalyst: Titanium compounds (e.g., TiCl₄) with organoaluminium compounds are used for the polymerization of alkenes.

The specific catalytic activity depends on the metal, its oxidation state, and the reaction conditions. Understanding these properties is crucial for designing efficient industrial processes.