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Transition metals—elements such as chromium, iron, and nickel—feature valence electrons in two shells, granting them unique catalytic properties. Because they can readily donate and accept electrons, they lower reaction barriers without altering thermodynamics.
Catalysts accelerate reactions by providing a lower-energy transition state, increasing the frequency of effective collisions while leaving reactants' identities unchanged. They do not shift the overall thermodynamic equilibrium.
While all transition metals belong to the d‑block, not every d‑block element qualifies. A true transition metal possesses an incompletely filled d‑orbital, such as scandium (Sc³⁺) or zinc (Zn²⁺), which lack d‑electrons and thus do not exhibit catalytic behavior.
Their ability to toggle between multiple oxidation states and form stable complexes enables them to exchange electrons with substrates efficiently. This versatility is the cornerstone of their catalytic performance.
Take copper as an example: it exists naturally as Cu⁺, Cu²⁺, and Cu³⁺. Its partially filled d‑orbitals allow it to act as either an electron donor or acceptor, facilitating redox processes essential to many catalytic cycles.
Transition metals coordinate with reactants to create metal complexes. If the transition state requires additional electrons, the metal undergoes oxidation or reduction to supply them. Conversely, it can sequester excess electron density, stabilizing intermediates and steering the reaction pathway.
Adsorption and absorption properties of both the metal and its complexes further influence catalytic efficiency, affecting how substrates approach and bind to the active site.
