Highly reactive metals normally occur in compounds rather than native form. Extraction may first convert an ore to a suitable oxide or solution. Metals below carbon in an appropriate reactivity comparison may be obtained by carbon or carbon-monoxide reduction of oxides. Very reactive metals often require electrolysis of molten compounds; water competes for reduction in aqueous solutions, so molten and aqueous processes can give different products.
An electrolytic cell uses external electrical energy to drive a non-spontaneous reaction. Its cathode is negative and carries out reduction; its anode is positive and carries out oxidation. Cations move toward the cathode and anions toward the anode. Electrons enter the cathode through the external supply and are withdrawn from the anode. Balance atoms and charge when writing each half-equation.
Example 1
Write cathode and anode half-equations for molten NaCl with inert electrodes.
Cathode: Na++e−→ Na. Anode: 2Cl−→ Cl2+2e−. Combine after doubling the cathode equation: 2NaCl(l)→2Na+ Cl2.
Example 2
Show conversion of ZnCO3 ore to ZnO, then reduction of the oxide by CO, assuming these steps are feasible.
Conversion: ZnCO3→ZnO+ CO2. Reduction: ZnO+ CO →Zn+ CO2. The first step prepares an oxide; in the second, zinc is reduced while CO is oxidised. Each equation balances atoms.
Watch out:Oxidation always occurs at the anode, but an electrolytic anode is positive while a galvanic anode is negative.
Metal extraction and electrolytic cells practice | Year 12 Chemistry | Ascent Arcade