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Chemistry: redox reactions
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Very detailed explanation of redox reactions that help students understand all aspects of this section. Simple and easy to understand.
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Redox ReactionsTransfer of electrons.
Change in charge.
Oxidation: loss of electrons
Reduction: gain of electrons
Oxidizing agents undergo reduction (left hand side of table.)
−¿ ¿
Eg: Fe2+ ¿+ 2e ¿ Fe – reduction half
Reducing agents undergo oxidation (right hand side of table.)
−¿¿
Eg: Mg Mg 2+¿+2 e ¿
Always two half reactions for redox reaction.
Net reaction = common factor for electrons.
At the anode: oxidation occurs (AN OX)
At the cathode: reduction occurs (RED CAT)
Anodes and cathodes are good electrical conductors.
Must have concentration of 1 mol . dm−3, pressure 1 atm and a temperature of 25 °C.
Spontaneous redox reaction
Exothermic reaction as energy is released.
Positive emf
Galvanic cells
Converts chemical energy to electrical energy.
Electrons flow through conducting wire attached to electrodes.
Is a battery.
Two half reactions that take place in separate containers.
Containers connected by a salt bridge.
Anode = negative electrode.
Cathode = positive electrode.
Salt bridge (contains saturated ionic solution of KNO3 Na2SO4)
Maintains electrical neutrality.
Completes the circuit.
Electrons never flow in the salt bridge.
Cell notation
Anode(s)/ion(aq)//ion(aq)/cathode(s)
Cell potential/emf
Measure of the ability of a cell to do work.
Measure of the difference in the reduction potential of two half-cells.
Comes from table.
Every half-cell potential value is relative to hydrogen’s cell potential – standard
hydrogen electrode (0,00V)
Eθ cell=Eθ cathode− Eθ anode (positive for galvanic)
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