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Introductory Chemistry , 3 rd Edition Nivaldo Tro

Introductory Chemistry , 3 rd Edition Nivaldo Tro. Chapter 16 Oxidation and Reduction. Roy Kennedy Massachusetts Bay Community College Wellesley Hills, MA. 2009, Prentice Hall. Oxidation – Reduction Reactions. Oxidation – reduction reactions are also called redox reactions .

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Introductory Chemistry , 3 rd Edition Nivaldo Tro

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  1. Introductory Chemistry, 3rd EditionNivaldo Tro Chapter 16 Oxidation and Reduction Roy Kennedy Massachusetts Bay Community College Wellesley Hills, MA 2009, Prentice Hall

  2. Oxidation–Reduction Reactions • Oxidation–reduction reactions are also called redox reactions. • All redox reactions involve the transfer of electrons from one atom to another. • Spontaneous redox reactions are generally exothermic, and we can use their released energy as a source of energy for other applications. • Convert the heat of combustion into mechanical energy to move our cars. • Use electrical energy in a car battery to start our car engine. Tro's Introductory Chemistry, Chapter 16

  3. Combustion Reactions • Combustion reactions are always exothermic. • In combustion reactions, O2 combines with all the elements in another reactant to make the products. 4 Fe(s) + 3 O2(g) → 2 Fe2O3(s) + energy CH4(g) + 2 O2(g) → CO2(g) + 2 H2O(g) + energy Tro's Introductory Chemistry, Chapter 16

  4. Reverse of Combustion Reactions • Since combustion reactions are exothermic, their reverse reactions are endothermic. • The reverse of a combustion reaction involves the production of O2. energy + 2 Fe2O3(s) → 4 Fe(s) + 3 O2(g) energy + CO2(g) + 2 H2O(g) → CH4(g) + 2 O2(g) • Reactions in which O2 is gained or lost are redox reactions. Tro's Introductory Chemistry, Chapter 16

  5. Oxidation and Reduction:One Definition • When an element attaches to an oxygen during the course of a reaction it is generally being oxidized. • In CH4(g) + 2 O2(g) → CO2(g) + 2 H2O(g), C is being oxidized in this reaction, but H is not. • When an element loses an attachment to oxygen during the course of a reaction, it is generally being reduced. • In 2 Fe2O3(s) → 4 Fe(s) + 3 O2(g), the Fe is being reduced. • One definition of redox is the gain or loss of O, but it is not the best. Tro's Introductory Chemistry, Chapter 16

  6. Another Oxidation–Reduction • Consider the following reactions: 4 Na(s) + O2(g) → 2 Na2O(s) 2 Na(s) + Cl2(g) → 2 NaCl(s) • The reaction involves a metal reacting with a nonmetal. • In addition, both reactions involve the conversion of free elements into ions. 4 Na(s) + O2(g) → 2 Na+2O–(s) 2 Na(s) + Cl2(g) → 2 Na+Cl–(s) Tro's Introductory Chemistry, Chapter 16

  7. Ger Leo Oxidation and Reduction:Another Definition • In order to convert a free element into an ion, the atoms must gain or lose electrons. • Of course, if one atom loses electrons, another must accept them. • Reactions where electrons are transferred from one atom to another are redox reactions. • Atoms that lose electrons are being oxidized, atoms that gain electrons are being reduced. 2 Na(s) + Cl2(g) → 2 Na+Cl–(s) Na → Na+ + 1 e– (oxidation) Cl2+ 2 e– → 2 Cl– (reduction) Tro's Introductory Chemistry, Chapter 16

  8. Practice—Identify the Element Being Oxidized and the Element Being Reduced. • 2 C(s) + O2(g) → 2 CO(g) • Mg(s) + Cl2(g) → MgCl2(s) • Mg(s) + Fe2+(aq) → Mg2+(aq) + Fe(s) Tro's Introductory Chemistry, Chapter 16

  9. Practice—Identify the Element Being Oxidized and the Element Being Reduced, Continued. • 2 C(s) + O2(g) → 2 CO(g) • Mg(s) + Cl2(g) → MgCl2(s) • Mg(s) + Fe2+(aq) → Mg2+(aq) + Fe(s) C is oxidized because it is gaining an attachment to O. O is reduced; there has to be reduction and it’s the only other element. 2+ − 0 0 Mg is oxidized because it is becoming a cation by losing electrons. Cl is reduced because it is becoming an anion by gaining electrons. Mg is oxidized because it is becoming a cation by losing electrons. Fe2+ is reduced because it is gaining electrons to become neutral. Tro's Introductory Chemistry, Chapter 16

  10. Oxidation–Reduction • Oxidation and reduction must occur simultaneously. • If an atom loses electrons, another atom must take them. • The reactant that reduces an element in another reactant is called the reducing agent. • The reducing agent contains the element that is oxidized. • The reactant that oxidizes an element in another reactant is called the oxidizing agent. • The oxidizing agent contains the element that is reduced. 2 Na(s) + Cl2(g) → 2 Na+Cl–(s) Na is oxidized, Cl is reduced. Na is the reducing agent, Cl2 is the oxidizing agent. Tro's Introductory Chemistry, Chapter 16

  11. Electron Bookkeeping • For reactions that are not metal + nonmetal, or do not involve O2, we need a method for determining how the electrons are transferred. • Chemists assign a number to each element in a reaction called an oxidation state that allows them to determine the electron flow in the reaction. • Although they look like them, oxidation states are not ion charges! • Oxidation states are imaginary charges assigned based on a set of rules. • Ion charges are real, measurable charges. Tro's Introductory Chemistry, Chapter 16

  12. Rules for Assigning Oxidation States, Continued • In their compounds, nonmetals have oxidation states according to the table below. • Nonmetals higher on the table take priority. Tro's Introductory Chemistry, Chapter 16

  13. Practice—Assign an Oxidation State to Each Element in the Following: • F2 • Mg2+ • KCl • SO2 • PO43− • BaO2 Tro's Introductory Chemistry, Chapter 16

  14. Practice—Assign an Oxidation State to Each Element in the Following, Continued: • F2F = 0 (Rule 1) • Mg2+ Mg = +2 (Rule 2) • KCl K = +1 (Rule 4a) and Cl = -1 (Rule 5) • SO2 O = -2 (Rule 5) and S = +4 (Rule 3a) • PO43−O = -2 (Rule 5) and P = +5 (Rule 3b) • BaOBa = +2 (Rule 4b) and O = -2 Tro's Introductory Chemistry, Chapter 16

  15. oxidation reduction Oxidation and Reduction:A Better Definition • Oxidation occurs when an atom’s oxidation state increases during a reaction. • Reduction occurs when an atom’s oxidation state decreases during a reaction. CH4 + 2 O2 → CO2 + 2 H2O -4 +1 0+4 –2 +1 -2 Tro's Introductory Chemistry, Chapter 16

  16. Practice—Assign Oxidation States and Identify the Oxidizing and Reducing Agents in Each of the Following: • 3 H2S + 2 NO3– + 2 H+® 3S + 2 NO + 4 H2O • MnO2 + 4 HBr ® MnBr2 + Br2 + 2 H2O Tro's Introductory Chemistry, Chapter 16

  17. oxidation reduction oxidation reduction Practice—Assign Oxidation States and Identify the Oxidizing and Reducing Agents in Each of the Following, Continued: reducing agent oxidizing agent • 3 H2S + 2 NO3– + 2 H+® 3S + 2 NO + 4 H2O • MnO2 + 4 HBr ® MnBr2 + Br2 + 2 H2O +1 -2 +5 -2 +1 0 +2 -2 +1 -2 Oxidizing agent reducing agent +4 -2 +1 -1 +2 -1 0 +1 -2 Tro's Introductory Chemistry, Chapter 16

  18. Tendency to Lose Electrons • Some metals have a greater tendency to lose electrons than others. • Metallic-free elements are always oxidized. • The greater the tendency of a metal to lose electrons, the easier it is to oxidize. • The greater the tendency of a metal to lose electrons, the harder it is to reduce its cations. • If Metal A has a greater tendency to lose electrons than Metal B, then: A(s) + B+(aq)  A+(aq) + B(s), but: A+(aq) + B(s)  no reaction. Tro's Introductory Chemistry, Chapter 16 42

  19. displace H2 displace H2 displace H2 displace H2 K Ba Sr Ca Na Mg Al Mn Zn Cr Fe Cd Co Ni Sn Pb H Sb As Bi Cu Hg Ag Pd Pt Au K Ba Sr Ca Na Mg Al Mn Zn Cr Fe Cd Co Ni Sn Pb H Sb As Bi Cu Hg Ag Pd Pt Au K Ba Sr Ca Na Mg Al Mn Zn Cr Fe Cd Co Ni Sn Pb H Sb As Bi Cu Hg Ag Pd Pt Au K Ba Sr Ca Na Mg Al Mn Zn Cr Fe Cd Co Ni Sn Pb H Sb As Bi Cu Hg Ag Pd Pt Au from cold H2O from cold H2O from cold H2O from cold H2O from steam from steam from steam from steam react with O2 in the air to make oxides react with O2 in the air to make oxides react with O2 in the air to make oxides react with O2 in the air to make oxides Gold is at the bottom, so it is very unreactive. from acids from acids from acids from acids Zn + 2 H+® H2 + Zn2+ Fe is above Cu, so Cu metal will not displace Fe2+ Fe is below Zn, so Zn metal will displace Fe2+. Zn is above H, so Zn will react with acids Activity Series of Metals • Listing of metals by reactivity. • Free metal higher on the list displaces metal cation lower on the list. • Metals above H will dissolve in acid: Zn + Fe2+® Fe + Zn2+ Cu + Fe2+® no reaction

  20. Mg will react with Cu2+ to form Mg2+ and Cu metal. Cu will not react with Mg2+. Mg is above Cu on the activity series. Tro's Introductory Chemistry, Chapter 16

  21. Table of Oxidation Half-Reactions 45

  22. Electrical Current • When we talk about the current of a liquid in a stream, we are discussing the amount of water that passes by in a given period of time. • When we discuss electric current, we are discussing the amount of electric charge that passes a point in a given period of time. • Whether as electrons flowing through a wire or ions flowing through a solution. Tro's Introductory Chemistry, Chapter 16 50

  23. Redox Reactions and Current • Redox reactions involve the transfer of electrons from one substance to another. • Therefore, redox reactions have the potential to generate an electric current. • In order to use that current, we need to separate the place where oxidation is occurring from the place that reduction is occurring. Tro's Introductory Chemistry, Chapter 16 51

  24. Electric Current Flowing Directly Between Atoms 52

  25. Electric Current Flowing Indirectly Between Atoms Tro's Introductory Chemistry, Chapter 16 53

  26. Electrochemical Cells • Electrochemistry is the study of redox reactions that produce or require an electric current. • The conversion between chemical energy and electrical energy is carried out in an electrochemical cell. • Spontaneous redox reactions take place in a voltaic cell. • Also known as galvanic cells. • Batteries are voltaic cells. • Nonspontaneous redox reactions can be made to occur in an electrolytic cell by the addition of electrical energy. Tro's Introductory Chemistry, Chapter 16

  27. Electrochemical Cells, Continued • Oxidation and reduction reactions kept separate. • Half-cells. • Electron flow through a wire, along with ion flow through a solution, constitutes an electric circuit. • Requires a conductive solid (metal or graphite) electrode to allow the transfer of electrons. • Through external circuit. • Ion exchange between the two halves of the system. • Electrolyte. Tro's Introductory Chemistry, Chapter 16

  28. Electrodes • Anode • Electrode where oxidation occurs. • Anions attracted to it. • Connected to positive end of battery in electrolytic cell. • Loses weight in electrolytic cell. • Cathode • Electrode where reduction occurs. • Cations attracted to it. • Connected to negative end of battery in electrolytic cell. • Gains weight in electrolytic cell. • Electrode where plating takes place in electroplating. Tro's Introductory Chemistry, Chapter 16

  29. Voltaic Cell

  30. Current and Voltage • The number of electrons that flow through the system per second is the current. • Electrode surface area dictates the number of electrons that can flow. • The amount of force pushing the electrons through the wire is the voltage. • The farther the metals are separated on the activity series, the larger the voltage will be. Tro's Introductory Chemistry, Chapter 16

  31. Current The number of electrons that pass a point each second is called the current of the electricity. The amount of water that passes a point each second is called the current of the river. Tro's Introductory Chemistry, Chapter 16

  32. Voltage Voltage is the force pushing the electrons down the wire. Gravity is the force pulling the water down the river. Tro's Introductory Chemistry, Chapter 16

  33. LeClanché’s Acidic Dry Cell • Electrolyte in paste form. • ZnCl2 + NH4Cl. • Or MgBr2. • Anode = Zn (or Mg). • Zn(s) ®Zn2+(aq) + 2 e- • Cathode = graphite rod. • MnO2 is reduced. 2 MnO2(s) + 2 NH4+(aq) + 2 H2O(l) + 2 e- ® 2 NH4OH(aq) + 2 Mn(O)OH(s) • Cell voltage = 1.5 v. • Expensive, nonrechargeable, heavy, easily corroded. Tro's Introductory Chemistry, Chapter 16

  34. Alkaline Dry Cell • Same basic cell as acidic dry cell, except electrolyte is alkaline KOH paste. • Anode = Zn (or Mg). Zn(s) ®Zn2+(aq) + 2 e- • Cathode = brass rod. • MnO2 is reduced. 2 MnO2(s) + 2 NH4+(aq) + 2 H2O(l) + 2 e- ® 2 NH4OH(aq) + 2 Mn(O)OH(s) • Cell voltage = 1.54 v. • Longer shelf life than acidic dry cells and rechargeable; little corrosion of zinc. Tro's Introductory Chemistry, Chapter 16

  35. Lead Storage Battery • Six cells in series. • Electrolyte = 6 M H2SO4. • Anode = Pb. Pb(s) + SO42-(aq) ®PbSO4(s) + 2 e- • Cathode = Pb coated with PbO2. • PbO2 is reduced. PbO2(s) + 4 H+(aq) + SO42-(aq) + 2 e- ® PbSO4(s) + 2 H2O(l) • Cell voltage = 2.09 v. • Rechargeable, heavy. Tro's Introductory Chemistry, Chapter 16

  36. Fuel Cells • Like batteries in which reactants are constantly being added. • So it never runs down! • Anode and cathode both Pt-coated metal. • Electrolyte is OH– solution. • Anode reaction: 2 H2 + 4 OH– → 4 H2O(l) + 4 e-. • Cathode reaction: O2 + 4 H2O + 4 e- → 4 OH–.

  37. Electrolysis • Electrolysis is the process of using electricity to break a compound apart. • Electrolysis is done in an electrolytic cell. • Electrolytic cells can be used to separate elements from their compounds. • Generate H2 from water for fuel cells. • Recover metals from their ores. Tro's Introductory Chemistry, Chapter 16

  38. Electrolytic Cell • The + terminal of the battery = anode. • The - terminal of the battery = cathode. • Cations attracted to the cathode; anions attracted to the anode. • Cations pick up electrons from the cathode and are reduced; anions release electrons to the anode and are oxidized. • In electroplating, the work piece is the cathode. • Cations are reduced at the cathode and plate onto the surface. • The anode is made of the plate metal, the anode oxidizes and replaces the metal cations lost from the solution. Tro's Introductory Chemistry, Chapter 16

  39. Electrolytic Cell—Electroplating Tro's Introductory Chemistry, Chapter 16

  40. Corrosion Corrosion is the spontaneous oxidation of a metal by chemicals in the environment. Since many materials we use are active metals, corrosion can be a very big problem. Tro's Introductory Chemistry, Chapter 16 70

  41. Preventing Corrosion One way to reduce or slow corrosion is to coat the metal surface to keep it from contacting corrosive chemicals in the environment. Paint. Some metals, like Al, form an oxide that strongly attaches to the metal surface, preventing the rest from corroding. Another method to protect one metal is to attach it to a more reactive metal that is cheap. Sacrificial electrode. Tro's Introductory Chemistry, Chapter 16 71

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