a.k.a Electrochemistry

1. Redox! a.k.a. Oxidation-Reduction a.k.a Electrochemistry

2. Review of Oxidation numbers The charge the atom would have in a molecule (or an ionic compound) if electrons were completely transferred. • Free elements (uncombined state) have an oxidation number of zero. Na, Be, K, Pb, H2, O2, P4 = 0 • In monatomic ions, the oxidation number is equal to the charge on the ion. Li+, Li =+1; Fe3+, Fe =+3; O2-, O =-2 • The oxidation number of oxygen is usually–2. In H2O2 it is–1. 4.4

3. Rules for Assigning Oxidation Numbers • The oxidation number of any uncombined element is zero. • The oxidation number of a monatomic ion equals its charge.

4. Rules for Assigning Oxidation Numbers • The oxidation number of oxygen in compounds is -2, except in peroxides, such as H2O2 where it is -1. • The oxidation number of hydrogen in compounds is +1, except in metal hydrides, like NaH, where it is -1.

5. Rules for Assigning Oxidation Numbers • The sum of the oxidation numbers of the atoms in the compound must equal 0. 2(+1) + (-2) = 0 H O (+2) + 2(-2) + 2(+1) = 0 Ca O H

6. Rules for Assigning Oxidation Numbers • The sum of the oxidation numbers in the formula of a polyatomic ion is equal to its ionic charge. X + 4(-2) = -2 S O X + 3(-2) = -1 N O  X = +6  X = +5

7. Electron transfer reactions are oxidation-reduction or redox reactions. Results in the generation of an electric current (electricity) or be caused by imposing an electric current. Therefore, this field of chemistry is often called ELECTROCHEMISTRY. Electron Transfer Reactions

8. 2Mg (s) + O2 (g) 2MgO (s) 2Mg 2Mg2+ + 4e- O2 + 4e- 2O2- • Electrochemical processes are oxidation-reduction reactions in which: • the energy released by a spontaneous reaction is converted to electricity or • electrical energy is used to cause a nonspontaneous reaction to occur 0 0 2+ 2- Oxidation half-reaction (lose e-) Reduction half-reaction (gain e-) 19.1

9. OXIDATION—loss of electron(s) by a species; increase in oxidation number; increase in oxygen. REDUCTION—gain of electron(s); decrease in oxidation number; decrease in oxygen; increase in hydrogen. OXIDIZING AGENT—electron acceptor; species is reduced. (an agent facilitates something; ex. Travel agents don’t travel, they facilitate travel) REDUCING AGENT—electron donor; species is oxidized. Terminology for Redox Reactions

10. You can’t have one… without the other! • Reduction (gaining electrons) can’t happen without an oxidation to provide the electrons. • You can’t have 2 oxidations or 2 reductions in the same equation. Reduction has to occur at the cost of oxidation LEO the lion says GER! ose lectrons xidation ain lectrons eduction GER!

11. Another way to remember • OIL RIG s s xidation ose eduction ain

12. Oxidation and Reduction (Redox) • Redox currently says that electrons are transferred between reactants • Mg + S → Mg2+ + S2- • The magnesium atom changes to a magnesium ion by losing 2 electrons, and is thus oxidized • The sulfur atom is changed to a sulfide ion by gaining 2 electrons, and is thus reduced.

13. Oxidation and Reduction (Redox) Each sodium atom loses one electron: Each chlorine atom gains one electron:

14. LEO says GER : Lose Electrons = Oxidation Sodium is oxidized Gain Electrons = Reduction Chlorine is reduced

15. LEO says GER : - Losing electrons is oxidation, and the substance that loses the electrons is called the reducing agent. - Gaining electrons is reduction, and the substance that gains the electrons is called the oxidizing agent. Mg(s) + S(s) → MgS(s) Mg is the reducing agent Mg is oxidized – loses e- S is the oxidizing agent S is reduced – gains e-

16. Not All Reactions are Redox Reactions - Reactions in which there has been no change in oxidation number are not redox reactions. Examples:

17. Reducing Agents and Oxidizing Agents • An increase in oxidation number = oxidation • A decrease in oxidation number = reduction Sodium is oxidized – it is the reducing agent Chlorine is reduced – it is the oxidizing agent

18. Trends in Oxidation and Reduction • Active metals: • Lose electrons easily • Are easily oxidized • Are strong reducing agents • Active nonmetals: • Gain electrons easily • Are easily reduced • Are strong oxidizing agents

19. Identifying Redox Equations • In general, all chemical reactions can be assigned to one of two classes: • oxidation-reduction, in which electrons are transferred: • Single-replacement, combination, decomposition, and combustion • this second class has no electron transfer, and includes all others: • Double-replacement and acid-base reactions

20. Identifying Redox Equations • In an electrical storm, oxygen and nitrogen react to form nitrogen monoxide: • N2(g) + O2(g)→ 2NO(g) • Is this a redox reaction? • If the oxidation number of an element in a reacting species changes, then that element has undergone either oxidation or reduction; therefore, the reaction as a whole must be a redox. YES!

21. Balancing Redox Equations • It is essential to write a correctly balanced equation that represents what happens in a chemical reaction • Fortunately, two systematic methods are available, and are based on the fact that the total electrons gained in reduction equals the total lost in oxidation. The two methods: • Use oxidation-number changes • Use half-reactions

22. Using half-reactions • A half-reaction is an equation showing just the oxidation or just the reduction that takes place • they are then balanced separately, and finally combined • Step 1: write unbalanced equation in ionic form • Step 2: write separate half-reaction equations for oxidation and reduction • Step 3: balance the atoms in the half-reactions

23. Using half-reactions • continued • Step 4: add enough electrons to one side of each half-reaction to balance the charges • Step 5: multiply each half-reaction by a number to make the electrons equal in both • Step 6: add the balanced half-reactions to show an overall equation • Step 7: add the spectator ions and balance the equation

24. To obtain a useful current, we separate the oxidizing and reducing agents so that electron transfer occurs thru an external wire. CHEMICAL CHANGE --->ELECTRIC CURRENT This is accomplished in a GALVANIC or VOLTAIC cell. A group of such cells is called a battery. http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/galvan5.swf

25. Galvanic Cells anode oxidation cathode reduction - + spontaneous redox reaction 19.2

26. Zn (s) + Cu2+(aq) Cu (s) + Zn2+(aq) Galvanic Cells • The difference in electrical potential between the anode and cathode is called: • cell voltage • electromotive force (emf) • cell potential Cell Diagram [Cu2+] = 1 M & [Zn2+] = 1 M Zn (s) | Zn2+ (1 M) || Cu2+ (1 M) | Cu (s) anode cathode 19.2

27. 2e- + 2H+ (1 M) H2 (1 atm) Zn (s) Zn2+ (1 M) + 2e- Zn (s) + 2H+ (1 M) Zn2+ + H2 (1 atm) Standard Electrode Potentials Zn (s) | Zn2+ (1 M) || H+ (1 M) | H2 (1 atm) | Pt (s) Anode (oxidation): Cathode (reduction): 19.3

28. 2e- + 2H+ (1 M) H2 (1 atm) Standard Electrode Potentials Standard reduction potential (E0) is the voltage associated with a reduction reaction at an electrode when all solutes are 1 M and all gases are at 1 atm. Reduction Reaction E0= 0 V Standard hydrogen electrode (SHE) 19.3

29. E0 is for the reaction as written • The more positive E0 the greater the tendency for the substance to be reduced • The half-cell reactions are reversible • The sign of E0changes when the reaction is reversed • Changing the stoichiometric coefficients of a half-cell reaction does not change the value of E0 19.3