General Trends in Properties of First-Row (Transition Metals)

The first-row transition metals, or 3d-series, are elements from Scandium (Z=21) to Zinc (Z=30). They are called "transition" elements because their properties are a transition between the highly reactive s-block metals and the covalent p-block elements. Their unique properties arise from the progressive filling of the 3d-orbitals.

1.0Transition Elements

According to IUPAC, a transition element is an element that has a partially filled d subshell in its ground state or in any of its common oxidation states.

• The d-block includes groups 3 through 12.
• Elements like Zinc (Zn), Cadmium (Cd), and Mercury (Hg) are conside red d-block elements but are not true transition elements because their d-orbitals are completely filled in their ground state and most common oxidation states  $(Z{n}^{2+},C{d}^{2+},H{g}^{2+})$

2.0First-Row Transition Metals

Elements in which 3d orbitals are progressively filled, spanning from Scandium (Sc) to Zinc (Zn).

Characteristic property: Can form one or more stable oxidation states, typically +2 and higher.

3.0Electronic Configuration of First-Row Transition Metals

The general electronic configuration of the first-row transition metals is $[Ar]3d^{1-10}4s^{1-2}$

Electrons are added to the 3d-orbital after the 4s-orbital is filled, as the 4s orbital has a lower energy than the 3d orbital.

There are two key exceptions to this trend due to the extra stability associated with half-filled (d5) and fully-filled (d10) d-orbitals.

• Chromium (Cr, Z=24): Expected configuration is $[Ar]3d^{4}4s^2$, but the observed configuration is $[Ar]3d^{5}4s^1$. This provides a stable, half-filled d-subshell.
• Copper (Cu, Z=29): Expected configuration is $[Ar]3d^{9}4s^2$, but the observed configuration is $[Ar]3d^{10}4s^1$. This provides a stable, fully-filled d-subshell.

4.0General Properties of Transition Elements

The presence of unpaired d-electrons and the small energy difference between the (n−1)d and ns orbitals give rise to several characteristic properties.

1. Atomic and Ionic Radii

The atomic radii of the first-row transition metals generally decrease gradually from left to right. This is because, as the atomic number increases, the effective nuclear charge increases, pulling the electrons closer to the nucleus. However, this decrease is less pronounced than in s- and p-block elements due to the increasing shielding effect of the d-electrons.

2. Ionization Enthalpy

The first ionization enthalpy generally increases from left to right across the series. This is due to the increase in effective nuclear charge, which makes it more difficult to remove an electron. However, there are some irregularities due to the stable half-filled or fully-filled d-orbitals (e.g., Mn and Zn).

3. Variable Oxidation States

The ability to exhibit multiple oxidation states is a key characteristic. This is because of the small energy difference between the 4s and 3d orbitals, allowing electrons from both shells to be involved in bond formation.

• The most common oxidation state is +2 (due to the loss of two 4s electrons).
• Manganese (Mn) shows the highest oxidation state of +7 , (in \ce{MnO4-})which corresponds to the loss of all its 4s and 3d electrons.

4. Formation of Coloured Ions

Most transition metal ions and their compounds are coloured. This is because of the presence of partially filled d-orbitals (d^1 to d^9) . When white light falls on a compound, the d-electrons absorb energy and undergo a d-d transition to a higher energy level, resulting in the transmission of the complementary colour.

5. Catalytic Properties

Many transition metals and their compounds act as excellent catalysts. This is attributed to their ability to:

• Exhibit variable oxidation states.
• Form intermediate complexes with reactants.
• Provide a suitable surface for the reaction to occur.
• Examples: Iron in the Haber process for ammonia synthesis, Vanadium pentoxide (\ce{V_2O_5}) in the Contact process for sulfuric acid production.

6. Magnetic Properties

Transition metal ions are either paramagnetic (attracted by a magnetic field) or diamagnetic (repelled by a magnetic field).

• The presence of unpaired d-electrons causes paramagnetism. The more unpaired electrons, the greater the paramagnetism.
• Diamagnetism occurs when all d-electrons are paired. Examples include $S{c}^{3+}(d^0)$ and ${\text{Zn}}^{2+}(d^{10})$