1. Oxidation and Reduction TOPIC 9

2. REDOX REACTIONS REDOX = reduction & oxidation O 2 (g) + 2 H 2 (g)  2 H 2 O( s ) O 2 (g) + 2 H 2 (g)  2 H 2 O( s )

3. REDOX REACTIONS REDOX = reduction & oxidation Corrosion of aluminum 2 Al(s) + 3 Cu 2+ (aq)  2 Al 3+ (aq) + 3 Cu(s)

4. In all reactions IF something has been oxidized then something has also been reduced REDOX REACTIONS Cu(s) + 2 Ag + (aq)  Cu 2+ (aq) + 2 Ag(s)

5. REDOX REACTIONS

6. Why Study Redox Reactions Manufacturing metals CorrosionCorrosion BatteriesBatteries FuelsFuels

7. Redox reactions are characterized by ELECTRON TRANSFER between an electron donor and electron acceptor. Transfer leads to— 1. increase in oxidation number of some element = OXIDATION 2.decrease in oxidation number of some element = REDUCTION REDOX REACTIONS

8. O I L R I G Oxidation It Loses (electrons) Reduction It Gains (electrons)

9. 2.In simple ions, oxidation # = charge on ion. -1 for Cl - +2 for Mg 2+ -1 for Cl - +2 for Mg 2+ 2.In simple ions, oxidation # = charge on ion. -1 for Cl - +2 for Mg 2+ -1 for Cl - +2 for Mg 2+ 1.Atoms in the elemental state have oxidation # = 0. Zn O 2 I 2 S 8 Zn O 2 I 2 S 8 1.Atoms in the elemental state have oxidation # = 0. Zn O 2 I 2 S 8 Zn O 2 I 2 S 8 OXIDATION NUMBERS The electric charge an element APPEARS to have when electrons are counted by some arbitrary rules:

10. 6. Some elements have oxidation states that vary in different compounds depending on the other in different compounds depending on the other elements present. 6. Some elements have oxidation states that vary in different compounds depending on the other in different compounds depending on the other elements present. 5. The usual oxidation # for an element is the same as the charge on its most common group ion. (table of exceptions on p.164 of IB book) 5. The usual oxidation # for an element is the same as the charge on its most common group ion. (table of exceptions on p.164 of IB book) 4. The oxidation # of all the atoms in a polyatomic ion 4. The oxidation # of all the atoms in a polyatomic ion Must add up to the charge on the ion. 4. The oxidation # of all the atoms in a polyatomic ion 4. The oxidation # of all the atoms in a polyatomic ion Must add up to the charge on the ion. 3. Oxidation #’s of all the atoms in neutral compound must add up to 0

11. 9.1.4 Recognizing a Redox Reaction Corrosion of aluminum 2 Al(s) + 3 Cu 2+ (aq)  2 Al 3+ (aq) + 3 Cu(s) Al(s)  Al 3+ (aq) + 3 e - Ox. # of Al increases as e - are donated by the metal. Ox. # of Al increases as e - are donated by the metal. Therefore, Al is OXIDIZED Therefore, Al is OXIDIZED

12. 9.1.4 Recognizing a Redox Reaction Corrosion of aluminum 2 Al(s) + 3 Cu 2+ (aq)  2 Al 3+ (aq) + 3 Cu(s) Cu 2+ (aq) + 2 e -  Cu(s) Ox. # of Cu decreases as e - are accepted by the ion. Ox. # of Cu decreases as e - are accepted by the ion. Therefore, Cu is REDUCED Therefore, Cu is REDUCED

13. 9.2.1 Recognizing a Redox Reaction Notice that the 2 half-reactions add up to give the overall reaction —if we use 2 mol of Al and 3 mol of Cu 2+. 2 Al(s)  2 Al 3+ (aq) + 6 e - 2 Al(s)  2 Al 3+ (aq) + 6 e - 3 Cu 2+ (aq) + 6 e -  3 Cu(s) 3 Cu 2+ (aq) + 6 e -  3 Cu(s)----------------------------------------------------------- 2 Al(s) + 3 Cu 2+ (aq)  2 Al 3+ (aq) + 3 Cu(s) Final eqn. is balanced for mass and charge.

14. Oxidizing and Reducing Agents If a substance has been oxidized, it is called the reducing agent because it reduced the charge on the other substance. If a substance has been reduced, it is called the oxidizing agent because it increased the charge on the other substance. in all reactions IF something has been oxidized then something has also been reduced Remember, in all reactions IF something has been oxidized then something has also been reduced

15. Common Oxidizing and Reducing Agents HNO 3 is an oxidizing agent 2 K + 2 H 2 O  2 KOH + H 2 Metals (Na, K, Mg, Fe) are reducing agents Metals (Cu) are reducing agents Cu + HNO 3  Cu 2+ + NO 2

16. 9.3 Reactivity  Oxidizing and Reducing agents are not all equal strength.  Metals are reducing agents  Some metals will be stronger than others depending on their relative tendencies to lose or gain electrons.  The REACTIVITY (IB term) series helps us predict if a reaction will occur or not.

17. 9.3 Reactivity Series examples Using the reactivity series, will these reactions occur? ZnCl 2 (aq) + 2Ag(s)  2AgCl(s) + Zn(s) 2FeCl 3 (aq) + 3Mg(s)  3MgCl 2 (aq) + 2Fe(s)

18. 9.4 Voltaic Cells A voltaic cell is an electrochemical cell used to convert chemical energy into electrical energy. Electrical energy is produced in a voltaic cell by a spontaneous redox reaction within the cell.

19. 9.4 Voltaic Cells A voltaic cell consists of two half-cells. A half-cell is one part of a voltaic cell in which either oxidation or reduction occurs. –A typical half-cell consists of a piece of metal immersed in a solution of its ions.

20. 9.4 Voltaic Cells The half-cells are connected by a salt bridge, which is a tube containing a strong electrolyte, often potassium sulfate (K 2 SO 4 ). A porous plate may be used instead of a salt bridge. The salt bridge or porous plate allows ions to pass from one half-cell to the other but prevents the solutions from mixing completely.

21. 9.4 Voltaic Cells Constructing a Voltaic Cell In this voltaic cell, the electrons generated from the oxidation of Zn to Zn 2+ flow through the external circuit (the wire) into the copper strip. Anode (–) Cathode (+) Salt bridge Cotton plugs ZnSO 4 solution CuSO 2 solution Zn(s) Zn 2+ (aq) + 2e – Cu 2+ (aq) + 2e – Cu(s) Wiree–e– e–e–

22. 9.4 Voltaic Cells The driving force of such a voltaic cell is the spontaneous redox reaction between zinc metal and copper ions in solution. Anode (–) Cathode (+) Salt bridge Cotton plugs ZnSO 4 solution CuSO 2 solution Zn(s) Zn 2+ (aq) + 2e – Cu 2+ (aq) + 2e – Cu(s) Wiree–e– e–e–

23. 9.4 Voltaic Cells The zinc and copper strips in this voltaic cell serve as the electrodes. Anode (–) Cathode (+) Salt bridge Cotton plugs ZnSO 4 solution CuSO 2 solution Zn(s) Zn 2+ (aq) + 2e – Cu 2+ (aq) + 2e – Cu(s) Wiree–e– e–e–

24. 9.4 Votaic Cells An electrode is a conductor in a circuit that carries electrons to or from a substance other than a metal. The electrode at which oxidation occurs is called the anode. – Electrons are produced at the anode. – The anode is labeled the negative electrode in a voltaic cell The electrode at which reduction occurs is called the cathode. – Electrons are consumed at the cathode in a voltaic cell. – The cathode is labeled the positive electrode.

25. 9.4 Voltaic Cells How a Voltaic Cell Works These steps actually occur at the same time. The electrochemical process that occurs in a zinc-copper voltaic cell can best be described in a number of steps.

26. 9.4 Voltaic Cells How a Voltaic Cell Works Step 1 Electrons are produced at the zinc strip according to the oxidation half-reaction: Zn(s) → Zn 2+ (aq) + 2e – Anode (–) Cathode (+) Salt bridge Cotton plugs ZnSO 4 solution CuSO 2 solution Zn(s) Zn 2+ (aq) + 2e – Cu 2+ (aq) + 2e – Cu(s) Wiree–e– e–e–

27. 9.4 Voltaic Cells How a Voltaic Cell Works The electrons leave the zinc anode and pass through the external circuit to the copper strip. Step 2 Anode (–) Cathode (+) Salt bridge Cotton plugs ZnSO 4 solution CuSO 2 solution Zn(s) Zn 2+ (aq) + 2e – Cu 2+ (aq) + 2e – Cu(s) Wiree–e– e–e– If a lightbulb is in the circuit, the electron flow will cause it to light.

28. 9.4 Voltaic Cells How a Voltaic Cell Works Step 3 Electrons interact with copper ions in solution. There, the following reduction half-reaction occurs: Cu 2+ (aq) + 2e – → Cu(s) Anode (–) Cathode (+) Salt bridge Cotton plugs ZnSO 4 solution CuSO 2 solution Zn(s) Zn 2+ (aq) + 2e – Cu 2+ (aq) + 2e – Cu(s) Wiree–e– e–e–

29. 9.4 Voltaic Cells How a Voltaic Cell Works Step 4 To complete the circuit, both positive and negative ions move through the aqueous solutions via the salt bridge. Anode (–) Cathode (+) Salt bridge Cotton plugs ZnSO 4 solution CuSO 2 solution Zn(s) Zn 2+ (aq) + 2e – Cu 2+ (aq) + 2e – Cu(s) Wiree–e– e–e–

30. Applications of voltaic cells Dry cells are voltaic cells in which the electrolyte is a paste. Positive button (+) Graphite rod (cathode) Moist paste of MnO 2, ZnCl 2, NH 4 Cl 2, H 2 O, and graphite powder Zinc (anode) Negative end cap (–)

31. Applications of voltaic cells A lead storage battery is a group of voltaic cells connected together. A 12-V car battery consists of six voltaic cells connected together.

32. 9.5 Electrolytic cells The process in which electrical energy is used to bring about a chemical change is called electrolysis. You are already familiar with some results of electrolysis, such as gold-plated jewelry, chrome- plated automobile parts, and silver- plated dishes. http://www.youtube.com/watc h?v=z7f7dQF2KLA

33. Electrolytic Cells The apparatus in which electrolysis is carried out is an electrolytic cell. An electrolytic cell is an electrochemical cell used to cause a chemical change through the application of electrical energy. An electrolytic cell uses electrical energy (direct current) to make a nonspontaneous redox reaction proceed to completion.

34. In both voltaic and electrolytic cells, electrons flow from the anode to the cathode in the external circuit. Anode (oxidation) Cathode (reduction) Energy Battery Voltaic Cell Electrolytic Cell e–e– e–e– e–e– e–e–

35. The key difference between voltaic and electrolytic cells is that in a voltaic cell, the flow of electrons is the result of a spontaneous redox reaction, whereas in an electrolytic cell, electrons are caused to flow by an outside power source, such as a battery. Anode (oxidation) Cathode (reduction) Energy Battery Voltaic CellElectrolytic Cell e–e– e–e– e–e– e–e–

36. In a voltaic cell, the anode is the negative electrode and the cathode is the positive electrode. In an electrolytic cell, the cathode is considered the negative electrode. Anode (oxidation) Cathode (reduction) Energy Battery Voltaic CellElectrolytic Cell e–e– e–e– e–e– e–e–

37. 9.5.4 Electrolysis of Molten Sodium Chloride Sodium and chlorine are produced through the electrolysis of pure molten sodium chloride, rather than an aqueous solution of NaCl. Chlorine gas is produced at the anode. Molten sodium collects at the cathode.

38. The overall equation is the sum of the two half-reactions: Electrolysis of Molten Sodium Chloride Reduction: 2Na + (l) + 2e – → 2Na(l) 2NaCl(l) → 2Na(l) + Cl 2 (g) Oxidation: 2Cl – (l) → Cl 2 (g) + 2e –

39. Electrolysis of Molten Sodium Chloride The electrolytic cell in which this commercial process is carried out is called the Downs cell. The cell operates at a temperature of 801°C so that the sodium chloride is maintained in the molten state.