Electrolytic Cells use an external power supply to force a non-spontaneous redox reaction to occur. Consider the following half-reactions and overall redox reaction: 2Cl Zn 2 2e Zn 2 2Cl Cl 2(g) 2e E 1.36 V Zn(s) E 0.76 V Cl 2(g) Zn(s) E 2.12 V To get an idea of what an electrolytic cell, in general is, we’ll start by looking at two half-reactions and the overall redox reaction we get by adding them. Consider the following half-reactions and overall redox reaction: 2Cl Zn 2 2e Zn 2 2Cl Cl 2(g) 2e E 1.36 V Zn(s) E 0.76 V Cl 2(g) Zn(s) E 2.12 V We’ll start with the half-reaction we get by reversing the reduction of chlorine on the table, to get the (click) oxidation of chloride ions Consider the following half-reactions and overall redox reaction: 2Cl Zn 2 2e Zn 2 2Cl Cl 2(g) 2e E 1.36 V Zn(s) E 0.76 V Cl 2(g) Zn(s) E 2.12 V Notice, because we reversed the equation, (click) the sign on the E naught is switched from positive to negative. Consider the following half-reactions and overall redox reaction: 2Cl Zn 2 2e Zn 2 2Cl Cl 2(g) 2e E 1.36 V Zn(s ) E 0.76 V Cl 2( g ) Zn(s ) E 2.12 V Next, we’ll add the half-reaction for the reduction of zinc ions, the way it is on the table, with an E naught value of negative 0.76 Volts. Consider the following half-reactions and overall redox reaction: 2Cl Zn 2 2e Zn 2 2Cl Cl 2(g ) 2e E 1.36 V Zn(s) E 0.76 V Cl 2(g) Zn(s) E 2.12 V We’ll add these two half-reactions to get the overall redox equation. Consider the following half-reactions and overall redox reaction: 2Cl Zn 2 2e Zn 2 2Cl Cl 2(g ) 2e E 1.36 V Zn(s) E 0.76 V Cl 2(g) Zn(s) E 2.12 V electrons gained = electrons lost Notice the electrons gained by the zinc ion, are (click) equal to the electrons lost by the chloride ions. Consider the following half-reactions and overall redox reaction: 2Cl Zn 2 2e Zn 2 2Cl Cl 2(g ) 2e E 1.36 V Zn(s) E 0.76 V Cl 2(g) Zn(s) E 2.12 V Therefore, electrons can be cancelled before adding the half-reactions. Consider the following half-reactions and overall redox reaction: 2Cl Zn 2 2e Zn 2 2Cl Cl 2(g) 2e E 1.36 V Zn(s) E 0.76 V Cl 2(g) Zn( s) E 2.12 V On the left side we have Zn2+ plus 2 Cl minus Consider the following half-reactions and overall redox reaction: 2Cl Zn 2 2e Zn 2 2Cl Cl 2(g) 2e E 1.36 V Zn( s) E 0.76 V Zn(s) Cl 2( g) E 2.12 V And On the right side we have Zn solid plus Cl2 gas Consider the following half-reactions and overall redox reaction: 2Cl Zn 2 2e Zn 2 2Cl Cl 2(g) 2e E 1.36 V Zn(s) E 0.76 V Zn(s) Cl 2(g) E 2.12 V To get the E naught value for the overall redox equation, (click) we add –1.36 volts and –0.76 volts, to give us (click) –2.12 volts Consider the following half-reactions and overall redox reaction: 2Cl Zn 2 2e Zn 2 2Cl Cl 2(g) 2e E 1.36 V Zn(s) E 0.76 V Zn(s) Cl 2(g) E 2.12 V This redox reaction is non-spontaneous A negative E naught value for the overall redox reaction, means that (click) this redox reaction is non-spontaneous as written. Consider the following half-reactions and overall redox reaction: 2Cl Zn 2 2e Zn 2 2Cl Cl 2(g) 2e E 1.36 V Zn(s) E 0.76 V Zn(s) Cl 2(g) E 2.12 V This redox reaction is non-spontaneous So that means, if we simply mixed Zn2+ and Cl–, we would (click) NOT get solid zinc and chlorine gas. Consider the following half-reactions and overall redox reaction: 2Cl Zn 2 2e Zn 2 2Cl Cl 2(g ) 2e E 1.36 V Zn(s) E 0.76 V Zn( s) Cl 2(g) E 2.12 V But we can Force this reaction to occur, using a process called electrolysis. Electrolysis uses an external power supply or battery to force a non-spontaneous redox reaction to occur. Zn 2 2Cl Zn(s) Cl 2(g) E 2.12 V Electrolysis uses an external power supply or battery to force a (click) nonspontaneous redox reaction to occur. Inert Carbon Electrode Inert Carbon Electrode Electrolysis takes place in an electrolytic cell. We’ll look at the simplest type of electrolytic cell. This is what we call a Type 1 Electrolytic cell. We’ll start off with (click) two inert carbon electrodes in a single container. Power Supply Inert Carbon Electrode + – Then we add a direct current power supply and wires. Inert Carbon Electrode Power Supply + – Neutral + e– e– + e– + e– + e– + + + + e– e– e– + e– + e– + e– + e– Neutral e– + e– + e– + e– + Before the power supply is connected, the electrodes are both neutral. Using simplified symbols, we’ll represent a few protons by + signs and a few electrons as “e”s with a negative charge. Because they are neutral, Power Supply + – Positive Neutral + e– e– + e– + e– + e– + + + + e– e– e– + e– + e– + e– + e– Negative Neutral e– + e– + e– + e– + When we close the switch and connect the power supply, it takes electrons from the electrode attached to the positive terminal and pumps them on to the electrode attached to the negative terminal Power Supply + – Positive Has a deficiency of electrons + + + + + e– e– + e– + + e– + e– e– –+ + e e– e– –+ + e e– e– –+ + e e– e–e–+ The electrode on the left now has a deficiency of electrons, giving it a net positive charge Power Supply + – Positive Has a deficiency of electrons + + + + + e– e– + e– + + e– + Negative e– e– –+ + e e– e– –+ + e e– e– –+ + e e– e–e–+ Has an excess of electrons And the electrode on the left now has an excess of electrons, giving it a net negative charge Power Supply Positive + – + ANODE + + + + e– e– + e– + + e– + e– e– –+ + e e– e– –+ + e e– e– –+ + e e– e–e–+ In an electrolytic cell, the electrode attached to the + terminal of the power supply is called (click) the Anode Power Supply Positive + – + ANODE Negative + + + + e– e– + e– + + e– + e– e– –+ + e e– e– –+ + e e– e– –+ + e e– e–e–+ CATHODE And the electrode attached to the negative terminal of the power supply is called (click) the Cathode. Power Supply Positive + – + ANODE A+ Negative + + + + e– e– + e– + + e– Just remember A plus for anode, + e– e– –+ + e e– e– –+ + e e– e– –+ + e e– e–e–+ CATHODE Power Supply Positive + – + ANODE A+ Negative + + + + e– e– + e– + + e– + e– e– –+ + e e– e– –+ + e e– e– –+ + e e– e–e–+ CATHODE C– And C minus for Cathode. Remember, this ONLY works for an Electrolytic cell, not an electrochemical cell. Power Supply + – ANODE + e– e– + e– e– e– e– – e – e e– – e – e e– – e – e – e e– – CATHODE – We’ll redistribute the electrons and stop showing the protons, for simplicity. Power Supply + – ANODE + e– e– – e – e e– – e – e e– – e – e – e e– e– + – e– e– e– ZnCl2(l) We’ll add some molten Zinc chloride to the container CATHODE – Power Supply + – ANODE + e– e– + e– Cl– e– e– Cl– Zn2+ Cl– Cl– Zn2+ Cl– Zn2+ Cl– Cl– e– – e – e e– – e – e e– – e – e – e e– – CATHODE – Cl– Zn2+ ZnCl2(l) A molten salt like zinc chloride consists of ions that are in constant random motion. Power Supply + – ANODE + e– e– + e– e– Cl– e– Zn2+ e– – e – e e– – e – e e– – e – e – e e– – Cl– ZnCl2(l) Now we’ll focus on one zinc ion and two chloride ions. CATHODE – Power Supply + – ANODE + e– e– + e– e– Cl– e– Zn2+ e– – e – e e– – e – e e– – e – e – e e– – CATHODE – Cl– ZnCl2(l) The positive zinc ions will be attracted to the negative cathode while the negative chloride ions will be attracted to the positive anode. Power Supply + – ANODE + e– e– + e– Cl– e– e– Zn2+ Cl– e– – e – e e– – e – e e– – e – e – e e– CATHODE – ZnCl2(l) Now, we’ll concentrate on the cathode and see what happens there. – Power Supply + –e ANODE + –– e– e– + e– Cl– e– e– Zn2+ Cl– ZnCl2(l) The zinc ion will gain 2 electrons from the cathode, e– – e – e e– – e – e e– – e – e – e e– – CATHODE – Power Supply + – ANODE + e– – e e– + e– Cl– e– e– e– Zn2+ – Zn e Cl– ZnCl2(l) And turn into a zinc atom e– e– – e– – e – e e– – e – e – e e– CATHODE – Power Supply Zn2+ + 2e– Zn(s) + – ANODE + e– – e e– + e– Cl– Zn e– e– Cl– ZnCl2(l) The equation for what just happened is Zn2+, e– e– – e– – e – e e– – e – e – e e– CATHODE – Power Supply Zn2+ + 2e– Zn(s) + – ANODE + e– – e e– + e– Cl– e– e– Zn Cl– ZnCl2(l) Plus 2 electrons e– e– – e– – e – e e– – e – e – e e– CATHODE – Power Supply Zn2+ + 2e– Zn(s) + – ANODE + e– – e e– + e– Cl– e– e– Zn Cl– ZnCl2(l) Gives Zinc solid e– e– – e– – e – e e– – e – e – e e– CATHODE – Power Supply Zn2+ + 2e– Zn(s) + – ANODE + Reduction e– – e e– + e– Cl– e– e– Zn Cl– ZnCl2(l) This is an example of reduction. e– e– – e– – e – e e– – e – e – e e– CATHODE – Power Supply Zn2+ + 2e– Zn(s) + – ANODE + Reduction e– – e e– + e– Cl– e– e– Zn Cl– e– e– CATHODE – e– – e – e e– – e – e – e e– ZnCl2(l) So we see that reduction of the cation Zn2+ occurs at the cathode – Power Supply Zn2+ + 2e– Zn(s) + – ANODE + Reduction e– – e e– + e– Cl– e– e– Zn Cl– e– e– – e– – e – e e– – e – e – e e– ZnCl2(l) Now we’ll have a look at the anode and see what happens. CATHODE – Power Supply Zn2+ + 2e– Zn(s) + – ee–– ANODE + e– + e– e– e– Reduction e– – e e– Cl Cl– Zn e– Cl– Cl e– e– – e– – e – e e– – e – e – e e– CATHODE – ZnCl2(l) Each chloride ion will lose one electron and change into a chlorine atom (Click half way through) Power Supply Zn2+ + 2e– Zn(s) + – ANODE + Reduction e– – e e– + e– e– e– Cl Zn e– Cl e– e– – e– – e – e e– – e – e – e e– CATHODE – ZnCl2(l) The 2 chlorine atoms then join to form a molecule of Cl2 or chlorine gas Power Supply 2Cl– Cl2(g) + 2e– + – Oxidation ANODE + Zn2+ + 2e– Zn(s) Reduction e– – e e– + e– e– e– Cl Zn e– Cl e– e– – e– – e – e e– – e – e – e e– CATHODE – ZnCl2(l) The process taking place on the anode can be summarized by the equation (click), 2 Cl minus Power Supply 2Cl– Cl2(g) + 2e– + – Oxidation ANODE + Zn2+ + 2e– Zn(s) Reduction e– – e e– + e– e– e– Cl Zn e– Cl ZnCl2(l) Gives Cl2 gas e– e– – e– – e – e e– – e – e – e e– CATHODE – Power Supply 2Cl– Cl2(g) + 2e– + – Oxidation ANODE + Zn2+ + 2e– Zn(s) Reduction e– – e e– + e– e– e– Cl Zn e– Cl ZnCl2(l) Plus 2 electrons. e– e– – e– – e – e e– – e – e – e e– CATHODE – Power Supply 2Cl– Cl2(g) + 2e– + – Oxidation ANODE + Zn2+ + 2e– Zn(s) Reduction e– – e e– + e– e– e– Cl Zn e– Cl e– e– – e– – e – e e– – e – e – e e– CATHODE – ZnCl2(l) The process of chloride ions losing electrons (click) is called oxidation Power Supply 2Cl– Cl2(g) + 2e– + – Oxidation ANODE + Zn2+ + 2e– Zn(s) Reduction e– – e e– + e– e– e– Cl Zn e– Cl e– e– – e– – e – e e– – e – e – e e– CATHODE – ZnCl2(l) And we see that oxidation of chloride ions to chlorine gas takes place at the anode. Power Supply 2Cl– Cl2(g) + 2e– + – Oxidation ANODE + Zn2+ + 2e– Zn(s) + Reduction CATHODE Zn2+ Zn – ZnCl2(l) So again zinc ions are reduced at the cathode to form zinc atoms. – Power Supply 2Cl– Cl2(g) + 2e– + – Oxidation ANODE + Zn2+ + 2e– Zn(s) + Reduction CATHODE Cl– Cl Zn Cl Cl– – ZnCl2(l) And chloride ions are oxidized at the anode to form chlorine gas, Cl2. – Power Supply 2Cl– Cl2(g) + 2e– + – Oxidation ANODE Zn2+ + 2e– Zn(s) Reduction e– CATHODE + Zn – 2Cl– Cl2(g) + 2e– – Zn2+ + 2e– Zn(s) ZnCl2(l) So as this cell operates, we can visualize zinc metal growing on the surface of the cathode and bubbles of chlorine gas forming on the anode 2Cl– Cl2(g) + 2e– Oxidation Power Supply Zn2+ + 2e– Zn(s) + – Reduction CATHODE ANODE + Zn – 2Cl– Cl2(g) + 2e– Product at Anode Zn2+ + 2e– Zn(s) ZnCl2(l) So we can say the product at the anode is chlorine gas – 2Cl– Cl2(g) + 2e– Oxidation Power Supply Zn2+ + 2e– Zn(s) + – Reduction CATHODE ANODE + Zn – 2Cl– Cl2(g) + 2e– Product at Anode Zn2+ + 2e– Zn(s) ZnCl2(l) And the product at the cathode is zinc metal. Product at Cathode – Molten zinc chloride is electrolyzed. ZnCl 2( l ) 2 Zn( l ) Reduction at Cathode 2Cl ( l ) Oxidation at Anode Zn 2 2e Zn(s) 2Cl Cl 2(g) 2e Product at Cathode is Zn(s) Overall Redox Reaction: Product at Anode is Cl2(g) Zn 2 2Cl Zn(s) Cl 2(g) Now that we’ve seen how this cell works, we’ll show you a process you can use for any questions involving the electrolysis of a molten salt. We’ll use our example here of molten zinc chloride. Molten zinc chloride is electrolyzed. ZnCl 2( l ) 2 Zn( l ) Reduction at Cathode 2Cl ( l ) Oxidation at Anode Zn 2 2e Zn(s) 2Cl Cl 2(g) 2e Product at Cathode is Zn(s) Overall Redox Reaction: Product at Anode is Cl2(g) Zn 2 2Cl Zn(s) Cl 2(g) Molten salts always consist of mobile positive and negative ions, so we write the dissociation equation for the salt forming liquid ions. We have ZnCl2 liquid Molten zinc chloride is electrolyzed. ZnCl 2( l ) 2 Zn( l ) Reduction at Cathode 2Cl ( l ) Oxidation at Anode Zn 2 2e Zn(s) 2Cl Cl 2(g) 2e Product at Cathode is Zn(s) Overall Redox Reaction: Forms Zn2+ liquid Product at Anode is Cl2(g) Zn 2 2Cl Zn(s) Cl 2(g) Molten zinc chloride is electrolyzed. ZnCl 2( l ) 2 Zn( l ) Reduction at Cathode 2Cl ( l ) Oxidation at Anode Zn 2 2e Zn(s) 2Cl Cl 2(g) 2e Product at Cathode is Zn(s) Overall Redox Reaction: Plus 2 Cl minus liquid Product at Anode is Cl2(g) Zn 2 2Cl Zn(s) Cl 2(g) ZnCl 2( l ) 2 Zn( l ) 2Cl ( l ) C– Reduction at Cathode Oxidation at Anode Zn 2 2e Zn(s) 2Cl Cl 2(g) 2e Product at Cathode is Zn(s) Overall Redox Reaction: Product at Anode is Cl2(g) Zn 2 2Cl Zn(s) Cl 2(g) We write C minus for cathode, underneath the positive ion. The positive ions are attracted to the negative cathode. ZnCl 2( l ) 2 Zn( l ) 2Cl ( l ) C– A+ Reduction at Cathode Oxidation at Anode Zn 2 2e Zn(s) 2Cl Cl 2(g) 2e Product at Cathode is Zn(s) Overall Redox Reaction: Product at Anode is Cl2(g) Zn 2 2Cl Zn(s) Cl 2(g) And we write A+ for anode, underneath the negative ions. The negative ions are attracted to the positive anode. ZnCl 2( l ) 2 Zn( l ) 2Cl ( l ) C– A+ Reduction at Cathode Oxidation at Anode Zn 2 2e Zn(s) 2Cl Cl 2(g) 2e Product at Cathode is Zn(s) Overall Redox Reaction: Reduction of Zn2+ ions occurs at the cathode. Product at Anode is Cl2(g) Zn 2 2Cl Zn(s) Cl 2(g) ZnCl 2( l ) 2 Zn( l ) 2Cl ( l ) C– A+ Reduction at Cathode Oxidation at Anode Zn 2 2e Zn(s) 2Cl Cl 2(g) 2e Product at Cathode is Zn(s) Overall Redox Reaction: Its half-reaction is: Zn2+ + 2e– Zn(s) . Product at Anode is Cl2(g) Zn 2 2Cl Zn(s) Cl 2(g) ZnCl 2( l ) 2 Zn( l ) 2Cl ( l ) C– A+ Reduction at Cathode Oxidation at Anode Zn 2 2e Zn(s) 2Cl Cl 2(g) 2e Product at Cathode is Zn(s) Overall Redox Reaction: And the product at the cathode is zinc solid Product at Anode is Cl2(g) Zn 2 2Cl Zn(s) Cl 2(g) ZnCl 2( l ) 2 Zn( l ) 2Cl ( l ) C– A+ Reduction at Cathode Oxidation at Anode Zn 2 2e Zn(s) 2Cl Cl 2(g) 2e Product at Cathode is Zn(s) Overall Redox Reaction: Oxidation of Cl– ions occurs at the anode Product at Anode is Cl2(g) Zn 2 2Cl Zn(s) Cl 2(g) ZnCl 2( l ) 2 Zn( l ) 2Cl ( l ) C– A+ Reduction at Cathode Oxidation at Anode Zn 2 2e Zn(s) 2Cl Cl 2(g) 2e Product at Cathode is Zn(s) Overall Redox Reaction: Its half-reaction is: 2Cl– Cl2(g) + 2e– . Product at Anode is Cl2(g) Zn 2 2Cl Zn(s) Cl 2(g) ZnCl 2( l ) 2 Zn( l ) 2Cl ( l ) C– A+ Reduction at Cathode Oxidation at Anode Zn 2 2e Zn(s) 2Cl Cl 2(g) 2e Product at Cathode is Zn(s) Overall Redox Reaction: And the product at the anode is chlorine gas Product at Anode is Cl2(g) Zn 2 2Cl Zn(s) Cl 2(g) ZnCl 2( l ) 2 Zn( l ) 2Cl ( l ) C– A+ Reduction at Cathode Oxidation at Anode Zn 2 2e Zn(s) 2Cl Cl 2(g) 2e Product at Cathode is Zn(s) Overall Redox Reaction: Product at Anode is Cl2(g) Zn 2 2Cl Zn(s) Cl 2(g) We get the equation for the overall redox reaction by adding up the two halfreactions. It is (click) Zn2+ + 2Cl– Zn(s) + Cl2(g) ZnCl 2( l ) 2 Zn( l ) 2Cl ( l ) C– A+ Reduction at Cathode Oxidation at Anode Zn 2 2e Zn(s) 2Cl Cl 2(g) 2e Product at Cathode is Zn(s) Overall Redox Reaction: Product at Anode is Cl2(g) Zn 2 2Cl Zn(s) Cl 2(g) Going through this process should greatly help you with any questions you get involving the electrolysis of molten salts, or what we call Type 1 electrolytic cells.