1. Chemical Reactions 2: Equilibrium & Oxidation-Reduction

2. Redox Reactions  Neutral atoms do not have charge since number of electrons equals number of protons (charge equals zero).  Charge is acquired when an atom gains (- charge) or loses (+ charge) electrons (last shell)

3. Redox Reactions Oxidation  Process of losing electrons (usually in last shell)  Most likely to occur to metals  Element “gains” charge (e.g. O 2- oxidizes to O, so charge changes from -2 to 0) (e.g. Zn oxidizes to Zn 2+, so charge changes from 0 to +2) Sodium lost one electron. It oxidized, so from Na to Na +

4. Redox Reactions Reduction  Process of gaining electrons (usually in last shell)  Most likely to occur to non-metals  Element “lose” charge (e.g. O oxidizes to O 2-, so charge changes from 0 to -2) (e.g. Cu 2+ oxidizes to Cu +, so charge changes from +2 to +1) Chlorine gained one electron. It reduced, so from Cl to Cl -

5. Redox Reactions  Oxidation half reaction produces electrons (M→M + + e - )  Reduction half reaction consumes electrons (N + e - →N - )

6. Redox Reactions Identify which reaction involves a reduction, and which an oxidation: _Zn → Zn 2+ + 2e - _S + 2e - → S 2- _Fe 2+ → Fe 3+ + e - _Al + 3e - → Al 3- Oxidation Reduction Oxidation Reduction

7. Redox Reactions Oxidizing agent: The one reactant that reduces in a redox reaction (N + e - →N - ) N reduces, so it is the oxidizing agent (makes M undergo oxidation) Reducing agent: The one reactant that oxidizes in a redox reaction (M→M + + e - ) M oxidizes, so it is the reducing agent (makes N undergo reduction)

8. Redox Reactions Copper. Cu 2+ (aq) + 2e - → Cu (s) Zinc. Zn (s) → Zn 2+ (aq) + 2e - Copper reduces. Zinc oxidizes Copper, oxidizing agent. Zinc, reducing agent

9. Redox Reactions  Oxidation and Reduction occur simultaneously  There cannot be one without the other  Both can be described by half-reactions  Total redox reactions needs to have same amount of electrons in both half reactions

10. Redox Reactions

13. Spontaneous Redox Reactions (Exothermic reactions) _Half-redox reactions are ranked according to their standard reduction potential, which is a measure of the tendency of an element to gain electrons _For a redox reaction to be spontaneous, the species acting as oxidizing agent (the one who reduces) must have a higher standard reduction potential than the species acting as reducing agent (the one who oxidizes)

14. Redox Reactions E° = -1.18V E° = -2.37V E° = 1.99V E° = -0.13V E° = -0.23V E° = -1.66V

15. Redox Reactions E° = -0.14V E° = -2.37V E° = 0.00V E° = -0.73V E° = 1.50V E° = 0.34V

16. Redox Reactions Volta’s cell was the first attempt to produce electricity. ***Even though Volta had little understanding of the way its cell worked, his discovery contributed to: _Development of electrochemistry _Discovery of new chemical elements

17. Redox Reactions Daniel’s cell _First truly usable cell _Very heavy and big equipment needed _Composed of: Anode (-) (electrode where oxidation takes place) Cathode (+) (electrode where reduction takes place) *Electrons flow from anode to cathode

18. Redox Reactions

19. Cell Potential Difference ΔE ° = E ° cathode - E ° anode (Reduction) (Oxidation) *each E ° is measured against the reduction potential of hydrogen electrode (zero)

20. Redox Reactions ΔE ° = E ° cathode - E ° anode ΔE ° = (0.34V) – (-0.76V) ΔE ° = 1.10 V Calculate ΔE ° if you replace Zn by Mg: ΔE ° = E ° cathode - E ° anode ΔE ° = (0.34V) –(-2.37V) ΔE ° = 2.71 V

21. Redox Reactions Cell notation Ag (s) /Ag + (aq) ||H + (aq) /H 2(g) Anode (Oxidation): Ag (s) → Ag + (aq) + 1e - Cathode (Reduction): 2H + (aq) + 2e - → H 2(g) Cell reaction: Ag (s) + 2H + (aq) → Ag + (aq) + H 2(g) || (Salt bridge): maintains electrical neutrality of solutions in half cells AnodeCathode Electrons move from anode to cathode

22. Redox Reactions