Transition Metals Overview

Transition Metals Overview

• Transition metals have an incomplete d-subshell or can form stable cations with incomplete d-subshell.
• Scandium and zinc are not considered transition metals despite being in the d-block.

Electronic Configuration

• The first five subshells can be abbreviated as [Ar], representing the argon core.
• Exceptions to the Aufbau Principle exist in the first row of d-block elements.
• Chromium has the electron configuration [Ar] 3d⁵ 4s¹.
• Copper has the electron configuration [Ar] 3d¹⁰ 4s¹.
• In ion formation, 4s electrons are removed before 3d electrons.

Properties of Transition Metals

• Transition metals can lose electrons to form positively charged ions, indicated by Roman numerals for oxidation states.
• Example: Iron (Fe) can form Fe²⁺ (Fe(II)) and Fe³⁺ (Fe(III)) ions.

Complex Ions

• Transition metals can form complex ions due to variable oxidation states.
• A complex ion consists of a central metal ion surrounded by ligands (molecules or ions).
• Ligands donate a pair of electrons, acting as Lewis bases (nucleophiles).
• Examples of iron(III) complexes include [Fe(NH₃)₆]³⁺, [Fe(OH)₆]³⁻, and [Fe(H₂O)₆]³⁺.

Color and Catalysis

• Compounds of transition metals are often colored.
• During catalysis, transition metals can change oxidation states, gaining or donating electrons.

Coordination Number and Ligand Types

• Coordination number refers to the number of coordinate bonds around a central metal atom or ion.
• Tetrahedral complexes (e.g., [CuCl₄]²⁻) and square planar complexes (e.g., [Ni(CN)₄]²⁻) have a coordination number of 4.
• Ligands can be multidentate, forming multiple dative bonds; for instance, EDTA⁴⁻ is a hexadentate ligand that forms 6 dative covalent bonds.