Electro Chemistry

1. Electro Chemistry Unit 5 Chemistry 12

2. Introduction • Electro chemistry is concerned with the conversion of chemical energy to electrical energy and visa versa. • Electrochemical reactions involve the movement of electrons;one reagent gains electrons, the other reagent looses electrons • Like acid/base reactions you can think of electrochemical reactions as a combination of 2 half reactions

3. Acids/Bases (In the rearview mirror) • Arrhenius: Acid donates H+ ions Base donates OH- ions • Bronstead-Lowery : Acid proton donor Base proton accepter • Lewis : Acid is an electron pair accepter Base an electron pair donor

4. Electro Chemical Reactions • 2 types (oxidation and reduction) • Oxidation 1/2 reaction in which the reacting species looses an electro • Reduction 1/2 reaction in which the reacting species gains an electron

5. Memory Trick “LEO the lion says GER” • LEO: loss of electrons = oxidation • GER: Gain of electrons = reduction We say Fe2+ is oxidized We say Cr3+was reduced

6. More Definitions • Oxidizing Agent: is the species that gets reduced (because it is what causes the other species to loose electrons) • Reducing Agent: is the species that gets oxidized (because it is what causes the other species to gain electrons)

7. Identification of oxidized/reduced species • Species that have been oxidized become more positive, because they have lost electrons (** the oxidized species is the reducing agent) • Species that have been reduced become more negative, because they have gained electrons. (** the reduced species is the oxidizing agent)

8. Predictions • Positive species will likely undergo reduction • Negative species will likely undergo oxidation • Neutral species will either undergo oxidation or reduction depending on their electron negativity.

9. Oxidation numbers • Oxidation number: is the charge that an atom would posses if the species containing the atom were made up of ions • Oxidation numbers are a useful way of deciding if a species has undergone oxidation or reduction • If the oxidation number of a species increases, then the species was oxidized • If the oxidation number of a species decreases, then the species was reduced.

10. Guidelines for AssigningOxidation Numbers • For ionic compounds (including acids/bases) the oxidation number is the charge on the ion • Group 1 have an oxidation # of plus 1 (hydrogen is sometimes an exception) • Group 2 have an oxidation # of plus 2. • Group 17 usually have an oxidation # of -1. • Group 16 usually have an oxidation # of -2. • THE SUM OF +VE and -VE CHARGES MUST BE EQUAL TO THE OVERALL CHARGE ON THE SPECIES.

11. Strengths of Reducing/oxidizing Agent • The relative strength of reducing/oxidizing agents are measured against the Standard Hydrogen 1/2 cell (often called the SHE electrode) • Standard potential tables are written from a reduction perspective

12. Spontaneity of Redox Reactions • For reagents to under go a redox reactions they must be on opposite sides of the std. reduction potential datasheet. (ie. One reagent must be able to under oxidation and the other reduction) • The reaction will be spontaneous if the reduction 1/2 cell is above the oxidation 1/2 cell in the reduction potential tables • The potential (voltage) of a spontaneous rx is always positive

13. Balancing RedoxReactions • Break the reaction into the two 1/2 cells • Balance each 1/2 cell • Use a stiochiometric ratio so that Total electrons gained=Total electrons lost • Add the two half reactions together canceling out species that appear on both sides of the reaction

14. Balancing 1/2 Cells (acid)MAJOR OH- • Balance Major ions first • Balance oxygen by adding H2O’s • Balance the hydrogen by adding H+ • Balance the charge by adding e-

15. Balancing 1/2 Cells (basic)MAJOR OH-(+water) Balance Major ions first Balance oxygen by adding H2O’s Balance the hydrogen by adding H+ Balance the charge by adding e- Use the dissociation of water to convert H+ into OH-

16. Preferred Redox Reactions • If more than one redox reaction is possible, the reaction which has the highest voltage potential will occur first. ( Farthest apart on the standard tables using “DISCO RULE”) • Example Rust protection using Zinc as a “sacrificial anode” • Zn -> Zn+2 +2e • Fe -> Fe+2 +2e • 1/2 O2(aq) + 2H+(10-7) +2e -> H2O

17. Galvanic/Voltaic/ElectrochemicalCells • An-Ox Ca-red An-OX Ca-Red

18. Electrolysis • Definition: is the process of supplying electrical energy to a molten ionic compound or ionic solution to produce a chemical change. Electrolysis supplies energy to non-spontaneous electro chemical reactions in which E0 is < O

19. Electrolytic or Electrolysis Cell • Electrolyte is a molten salt or ionic solution • For Electrolytic cells, The Redox reaction with the least negative voltage will occur first

20. Electrowinning • Electrowinning: Is an electrolytic cell that causes the electro deposition of metal from metallic ore/ liquid leach solution. The metal is deposited on the cathode. Usually oxygen is produced at the anode.

21. Electro-plating • Electro-plating is a process in which a metallic ion is reduced or “plated out” onto a cathode of a different material • The Electrolyte contains the metallic ion to be plated out. • The anode can be an inert electrode or the same material as the cathode

22. Electrorefining • Electrorefining is an electrolytic process in which impure metal is purified by oxidizing the impure metal at the anode and reduce the metallic ion produced at a cathode this process is voltage regulated such that only the reduction of the metal being purified plates out

23. Gottchas • Non-standard potentials • Aqeous solutions vs. molten salts