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Chapter 11 Oxidation ( ) and Reduction ( ) Acid-base reaction: Transfer of proton Oxidation and reduction: transfer of electron.

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Presentation on theme: "Chapter 11 Oxidation ( ) and Reduction ( ) Acid-base reaction: Transfer of proton Oxidation and reduction: transfer of electron."— Presentation transcript:

1 Chapter 11 Oxidation ( ) and Reduction ( ) Acid-base reaction: Transfer of proton Oxidation and reduction: transfer of electron

2 11.1 Oxidation is the loss of electrons and reduction is the gain of electrons Oxidation is the process whereby a reactant loses one or more electrons. Reduction is the opposite process whereby a reactant gains one or more electrons. Oxidation and reduction are complementary and simultaneous processes. 2Na + Cl 2 2NaCl 2Na 2Na+ + 2e - Oxidation Cl 2 + 2e - 2Cl - Reduction In this reaction, sodium is acting as a reducing agent ( ), which is any reactant that causes another reactant to be reduced. Conversely, the chlorine causes oxidation of the sodium and so it is acting as an oxidizing agent. Fig 11.1 formation of sodium chloride

3 Different elements have different oxidation and reduction tendencies Little tendency to lose or gain electrons Tendency to gain electrons Tendency to lose electrons

4 11.2 Photography works by selective oxidation and reduction

5 11.3 The energy of flowing electrons can be harnessed Electrochemistry ( ) is the study of the relationship between electrical energy and chemical change: Use oxidation- reduction reaction to produce an electric current or use an electric current to produce an oxidation-reduction reaction. Fig 11.7 the salt bridge completes the electric circuit

6 The electricity of a battery comes from oxidation- reduction reactions A voltaic cell, which is an all-in-one, self- contained unit, is called a battery. Batteries are either disposable or rechargeable. Principle of batteries: Two materials that oxidize and reduce each other are connected by a medium through which ions travel to balance an external flow of electrons.

7 Disposable batteries ZnCl 2 (ag) + 2NH 3 (g) Zn(NH 3 ) 2 Cl 2 (s) 2MnO 2 (s) + H 2 (g) Mn 2 O 3 (s) + H 2 O (l) Electrode ( ) cathode ( ): where chemicals are reduced. Anode ( ): where chemicals are oxidized. Fig 11.8 a common dry-cell battery with a graphite rod immersed in a paste of ammonium chloride, manganese dioxide, and zinc chloride Reduction 2NH e - 2NH 3 +H 2 Oxidation Zn Zn 2+ +2e -

8 Alkaline battery Zn (s) + 2 OH- (aq) ZnO (s) + H 2 O (l) +2 e- Oxidation 2MnO 2 (s) + 2e- Mn 2 O 3 (s) + 2OH- (aq) Reduction Fig11.19 Alkaline batteries last a lot longer than dry-cell batteries and give a steadier voltage, but they are expensive

9 Rechargeable battery engine Oxidation Pb + SO 4 2- PbSO 4 + 2e Reduction of elemental Pb to pb 2+ Reduction PbO 2 + SO H + + 2e PbSO 4 + 2H 2 O Oxidation of elemental pb 4+ to pb 2+ engine alternator Oxidation PbSO 4 + 2e Pb + SO 4 2- Reduction of elemental Pb 2+ to pb Reduction PbSO pb H 2 O PbO 2 + SO H + + 2e Oxidation of elemental pb 2+ to pb 4 + Fig (a) electrical energy from the battery forces the starter motor to start the engine. (b) the combustion of fuel keeps the engine running and provides energy to spin the alternator, which recharges the battery. Note that the battery has a reversed cathode- anode orientation during recharging

10 Fuel cells ( ) are highly efficient sources of electrical energy Oxidation 2H 2 +4OH - 4H 2 O+4e - Reduction 4e - +O 2 +2H 2 O 4OH - Porous graphite electrodes Fig11.11 the hydrogen- oxygen fuel cell

11 Fig because this bus is powered by a fuel cell, its tail pipe emits mostly water vapor

12 Electrical energy can produce chemical change Electrolysis ( ) is the use of electrical energy to produce chemical change. Electrical energy + 2H 2 O 2H 2 (g) + O 2 Fig the electrolysis of water produces hydrogen gas and oxygen gas in a 2:1ratio by volume, which is in accordance with the chemical formula for water:H2O. For this process to work, ions must be dissolved in the water so that the electricity can be conducted between the electrodes

13 Oxidation 2AlOF F - +C 2AlF CO 2 +4e - Reduction AlF e - Al +6F - Figure The melting point of aluminum oxide(2030 ) is too high for it to be efficiently electrolyzed to aluminum metal. When the oxide is mixed with the mineral cryolite. The melting point of the oxide drops to a more reasonable 980. A strong electric current passed through the molten aluminum oxide- cryolite mixture generates aluminum metal at the cathode, where aluminum ions pich up electrons and so are reduced to elemental aluminum Molten Al 2 O 3 +Na 3 AlF 6 mixture Power source cathode

14 11.4 Oxygen is responsible for corrosion and combustion Iron to rust: 4Fe + 3O 2 +3H 2 O 2Fe 2 O 3.3H 2 O A thin layer of protective layer was formed during the oxidation of aluminum. Since aluminum has higher tendency to be oxidized, a thin layer of aluminum on iron can protect iron from corresion. Cathodic protection Electroplating Figure As electrons flow into the hubcap and give it a negative charge, positively charged chromiun ions move from the solution to the hubcap and are reduced to chromium metal, which deposits as a coating on the hubcap. The solution is kept supplied with ions as chromium atoms in the cathode are oxidized to Cr 2+ ions

15 Combustion is also an oxidation-reduction reaction. Many well designed oxidation-reduction reactions happen in our bodies.

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