Compound examples of Cr, Mn, Fe, Co, Ni, Cu
Atomic structure and the periodic table • Properties of transition metals
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Definition of a transition metal and complexes
A transition metal is an element whose atoms form one or more stable ions with incompletely filled d-orbitals. Complexes form when transition-metal ions bind ligands (neutral molecules or anions) through coordinate covalent bonds, causing d-orbital splitting and characteristic properties. Limiting factor: some elements near the border of the d-block (e.g., zinc) do not show variable d-electron counts and so do not behave like transition metals.
Variable oxidation states and redox behaviour
Variable oxidation states arise because d-electrons participate in bonding and redox processes. Example: chromium forms Cr2O7 2− (Cr(VI), strongly oxidising) and Cr3+ (Cr(III), more stable); manganese forms MnO4 − (Mn(VII), strong oxidant) and Mn2+ (Mn(II), more stable). Cause → effect: higher oxidation states tend to be stronger oxidising agents and often show different colours and reactivity compared with lower states. Limiting factor: stability of an oxidation state depends on the ligand, pH and complex environment.
Colours caused by d–d transitions and charge transfer
Coloured ions result from absorption of visible light by electrons promoted between split d-orbitals (d–d transitions) or by charge-transfer transitions between ligand and metal. Example colours: [Cu(H2O)6]2+ is blue, [MnO4]− is deep purple, [Cr2O7]2− is orange. Cause → effect: larger ligand-field splitting shifts absorbed wavelength, altering observed colour. Limiting factor: ions with no available d–d transitions (d10) often appear colourless or white unless charge transfer occurs.
Paramagnetism and unpaired electrons
Paramagnetism arises from unpaired d-electrons in transition-metal ions. Example: Mn2+ (d5, high-spin) shows strong paramagnetism; Fe3+ (d5) is also strongly paramagnetic. Cause → effect: more unpaired electrons produce a stronger magnetic response and higher measured magnetic moment. Limiting factor: strong-field ligands can pair electrons and reduce paramagnetism.
Catalytic behaviour of transition-metal compounds
Transition-metal compounds catalyse reactions by providing alternative pathways with lower activation energy via temporary changes in oxidation state or ligand substitution. Example: iron compounds participate in redox cycles in catalytic processes; copper and cobalt compounds act as oxidation catalysts in laboratory reactions. Cause → effect: accessible multiple oxidation states facilitate catalytic cycles. Limiting factor: catalytic activity depends on the metal oxidation states, ligand environment and reaction conditions.
Toxicity and environmental reactivity
Some high oxidation state compounds, such as Cr(VI) and Mn(VII), act as strong oxidants and pose health and environmental hazards. Cause → effect: strong oxidising ability causes damage to organic material and biological systems. Limiting factor: hazard depends on concentration, solubility and speciation (e.g., chromate versus chromium(III) salts).
Specific compound examples and what they show
Cr compounds: dichromate (Cr2O7 2−, orange, strong oxidant) and Cr3+ salts (green, stable) show variable oxidation and colour differences. Mn compounds: permanganate (MnO4 −, purple, strong oxidant) and Mn2+ (pale pink) show redox extremes. Fe compounds: Fe2+ (pale green) and Fe3+ (yellow/brown) salts show redox interconversion and formation of rust (Fe2O3·xH2O). Co compounds: CoCl2 shows blue anhydrous and pink hydrated forms, illustrating coordination and hydration effects. Ni compounds: Ni2+ salts give green solutions and form complexes. Cu compounds: CuSO4·5H2O (blue) and [Cu(NH3)4(H2O)2]2+ (deep blue) show ligand exchange and characteristic blue colour.
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