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Oxidation and Reduction

Oxidation and Reduction. Topic 9. REDOX REACTIONS. REDOX = reduction & oxidation O 2 (g) + 2 H 2 (g)  2 H 2 O( s ). REDOX REACTIONS. REDOX = reduction & oxidation Corrosion of aluminum. 2 Al(s) + 3 Cu 2+ ( aq )  2 Al 3+ ( aq ) + 3 Cu(s). REDOX REACTIONS.

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Oxidation and Reduction

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  1. Oxidation and Reduction Topic 9

  2. REDOX REACTIONS REDOX = reduction & oxidation O2(g) + 2 H2(g) 2 H2O(s)

  3. REDOX REACTIONS REDOX = reduction & oxidation Corrosion of aluminum 2 Al(s) + 3 Cu2+(aq)  2 Al3+(aq) + 3 Cu(s)

  4. REDOX REACTIONS Cu(s) + 2 Ag+(aq)  Cu2+(aq) + 2 Ag(s) In all reactions IF something has been oxidized then something has also been reduced

  5. REDOX REACTIONS Cu(s) + 2 Ag+(aq)  Cu2+(aq) + 2 Ag(s)

  6. Fuels Why Study Redox Reactions Batteries Corrosion Manufacturing metals

  7. REDOX REACTIONS Redox reactions are characterized byELECTRON TRANSFER between an electron donor and electron acceptor. Transfer leads to— 1. increase in oxidation number of some element = OXIDATION • decrease in oxidation number of some element = REDUCTION O I L R I G

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

  9. OXIDATION NUMBERS The electric charge an element APPEARS to have when electrons are counted by some arbitrary rules: 1. Atoms in the elemental state have oxidation # = 0. Zn O2 I2 S8 2. In simple ions, oxidation # = charge on ion. -1 for Cl- +2 for Mg2+

  10. 3. Oxidation #’s of all the atoms in neutral compound must add up to 0 4. The oxidation # of all the atoms in a polyatomic ion Must add up to the charge on the ion. 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) 6. Some elements have oxidation states that vary in different compounds depending on the other elements present.

  11. 9.1.4 Recognizing a Redox Reaction Corrosion of aluminum 2 Al(s) + 3 Cu2+(aq)  2 Al3+(aq) + 3 Cu(s) Al(s)  Al3+(aq) + 3 e- • Ox. # of Al increases as e- are donated by the metal. • Therefore, Al is OXIDIZED

  12. 9.1.4 Recognizing a Redox Reaction Corrosion of aluminum 2 Al(s) + 3 Cu2+(aq)  2 Al3+(aq) + 3 Cu(s) Cu2+(aq) + 2 e-  Cu(s) • Ox. # of Cu decreases as e- are accepted by the ion. • 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 Cu2+. 2 Al(s) 2 Al3+(aq) + 6 e- 3 Cu2+(aq) + 6 e- 3 Cu(s) ----------------------------------------------------------- 2 Al(s) + 3 Cu2+(aq) 2 Al3+(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. • Remember, in all reactions IF something has been oxidized then something has also been reduced

  15. Metals (Cu) are reducing agents Metals (Na, K, Mg, Fe) are reducing agents HNO3 is an oxidizing agent Common Oxidizing and Reducing Agents Cu + HNO3Cu2+ + NO2 2 K + 2 H2O 2 KOH + H2

  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? • ZnCl2 (aq) + 2Ag(s)  2AgCl(s) + Zn(s) • 2FeCl3(aq) + 3Mg(s)  3MgCl2(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 (K2SO4). • 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. e– e– Wire Anode (–) Cathode (+) Salt bridge Cotton plugs ZnSO4 solution CuSO2 solution Zn(s) Zn2+(aq) + 2e– Cu2+(aq) + 2e– Cu(s) 9.4 Voltaic Cells • Constructing a Voltaic Cell In this voltaic cell, the electrons generated from the oxidation of Zn to Zn2+ flow through the external circuit (the wire) into the copper strip.

  22. e– e– Wire Anode (–) Cathode (+) Salt bridge Cotton plugs ZnSO4 solution CuSO2 solution Zn(s) Zn2+(aq) + 2e– Cu2+(aq) + 2e– Cu(s) 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.

  23. e– e– Wire Anode (–) Cathode (+) Salt bridge Cotton plugs ZnSO4 solution CuSO2 solution Zn(s) Zn2+(aq) + 2e– Cu2+(aq) + 2e– Cu(s) 9.4 Voltaic Cells The zinc and copper strips in this voltaic cell serve as the electrodes.

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

  26. e– e– Wire Anode (–) Cathode (+) Salt bridge Cotton plugs ZnSO4 solution CuSO2 solution Zn(s) Zn2+(aq) + 2e– Cu2+(aq) + 2e– Cu(s) 9.4 Voltaic Cells Step 1 • How a Voltaic Cell Works Electrons are produced at the zinc strip according to the oxidation half-reaction: Zn(s) → Zn2+(aq) + 2e–

  27. e– e– Wire Anode (–) Cathode (+) Salt bridge Cotton plugs ZnSO4 solution CuSO2 solution Zn(s) Zn2+(aq) + 2e– Cu2+(aq) + 2e– Cu(s) 9.4 Voltaic Cells If a lightbulb is in the circuit, the electron flow will cause it to light. Step 2 • How a Voltaic Cell Works The electrons leave the zinc anode and pass through the external circuit to the copper strip.

  28. e– e– Wire Anode (–) Cathode (+) Salt bridge Cotton plugs ZnSO4 solution CuSO2 solution Zn(s) Zn2+(aq) + 2e– Cu2+(aq) + 2e– Cu(s) 9.4 Voltaic Cells Step 3 • How a Voltaic Cell Works Electrons interact with copper ions in solution. There, the following reduction half-reaction occurs: Cu2+(aq) + 2e–→ Cu(s)

  29. e– e– Wire Anode (–) Cathode (+) Salt bridge Cotton plugs ZnSO4 solution CuSO2 solution Zn(s) Zn2+(aq) + 2e– Cu2+(aq) + 2e– Cu(s) 9.4 Voltaic Cells Step 4 • How a Voltaic Cell Works To complete the circuit, both positive and negative ions move through the aqueous solutions via the salt bridge.

  30. Positive button (+) Graphite rod (cathode) Moist paste of MnO2, ZnCl2, NH4Cl2, H2O, and graphite powder Zinc (anode) Negative end cap (–) Applications of voltaic cells • Dry cells are voltaic cells in which the electrolyte is a paste.

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

  35. Voltaic Cell Electrolytic Cell Battery e– e– e– e– Cathode (reduction) Anode (oxidation) Energy Anode (oxidation) Cathode (reduction) Energy 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.

  36. Voltaic Cell Electrolytic Cell Battery e– e– e– e– Cathode (reduction) Anode (oxidation) Energy Anode (oxidation) Cathode (reduction) Energy • 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.

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

  38. Oxidation: 2Cl–(l) → Cl2(g) + 2e– Reduction:2Na+(l) + 2e– → 2Na(l) 2NaCl(l) → 2Na(l) + Cl2(g) • Electrolysis of Molten Sodium Chloride • The overall equation is the sum of the two half-reactions:

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

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