Chemical Reactions 2: Equilibrium & Oxidation-Reduction

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. O2- oxidizes to O, so charge changes from -2 to 0) (e.g. Zn oxidizes to Zn2+, 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 O2-, so charge changes from 0 to -2) (e.g. Cu2+ 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 → Zn2+ + 2e- _S + 2e- → S2- _Fe2+→ Fe3++ e- _Al + 3e- → Al3- 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. Cu2+(aq) + 2e-→ Cu(s) Zinc. Zn(s)→ Zn2+(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

11. Redox Reactions

12. Redox Reactions

13. Redox Reactions 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 higherstandard 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. Redox Reactions 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)/H2(g) Anode (Oxidation): Ag(s)→Ag+(aq) +1e- Cathode (Reduction): 2H+(aq) +2e- →H2(g) Cell reaction: Ag(s) +2H+(aq) →Ag+(aq) +H2(g) || (Salt bridge): maintains electrical neutrality of solutions in half cells Anode Cathode Electrons move from anode to cathode

22. Redox Reactions

23. Redox Reactions

24. Redox Reactions