1 / 126

Chapter 20 Electrochemistry

Learn about electrochemical reactions, oxidation numbers, and oxidation-reduction processes in chemistry. Understand how to assign oxidation numbers and balance oxidation-reduction equations using the half-reaction method.

wilcox
Télécharger la présentation

Chapter 20 Electrochemistry

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Chemistry, The Central Science, 10th edition Theodore L. Brown; H. Eugene LeMay, Jr.; and Bruce E. Bursten Chapter 20Electrochemistry John D. Bookstaver St. Charles Community College St. Peters, MO  2006, Prentice Hall, Inc.

  2. Electrochemical Reactions In electrochemical reactions, electrons are transferred from one species to another.

  3. Oxidation Numbers In order to keep track of what loses electrons and what gains them, we assign oxidation numbers.

  4. Oxidation and Reduction • A species is oxidized when it loses electrons. • Here, zinc loses two electrons to go from neutral zinc metal to the Zn2+ ion.

  5. Oxidation and Reduction • A species is reduced when it gains electrons. • Here, each of the H+ gains an electron and they combine to form H2.

  6. Given the following reaction, what is chemical process is happening to Mg? Mg + 2H+ Mg2+ + H2 • oxidation • reduction • combustion • None of the above

  7. Given the following reaction, what chemical process is happening to H+? Mg + 2H+ Mg2+ + H2 • oxidation • reduction • combustion • None of the above

  8. Oxidation and Reduction • What is reduced is the oxidizing agent. • H+ oxidizes Zn by taking electrons from it. • What is oxidized is the reducing agent. • Zn reduces H+ by giving it electrons.

  9. Given the following reaction, which species is acting as the reducing agent? Mg + 2H+ Mg2+ + H2 • H+ • Mg • Mg2+ • H2

  10. Given the following reaction, which species is acting as the oxidizing agent? Mg + 2H+ Mg2+ + H2 • H2 • Mg • Mg2+ • H+

  11. Assigning Oxidation Numbers • Elements in their elemental form have an oxidation number of 0. • The oxidation number of a monatomic ion is the same as its charge.

  12. What is the oxidation state of magnesium metal? Mg + 2H+ Mg2+ + H2 • -2 • +1 • 0 • +2

  13. What is the oxidation state of hydrogen gas? Mg + 2H+ Mg2+ + H2 • -2 • +1 • 0 • +2

  14. Assigning Oxidation Numbers • Nonmetals tend to have negative oxidation numbers, although some are positive in certain compounds or ions. • Oxygen has an oxidation number of −2, except in the peroxide ion in which it has an oxidation number of −1. • Hydrogen is −1 when bonded to a metal, +1 when bonded to a nonmetal.

  15. What is the oxidation state of oxygen in the permanganate ion? MnO4-1 • -8 • -1 • -2 • -4

  16. What is the oxidation state of hydrogen in H2SO4? • +1 • -1 • +2 • None of the above

  17. What is the oxidation state of hydrogen in LiH? • +1 • -1 • +2 • None of the above

  18. Assigning Oxidation Numbers • Nonmetals tend to have negative oxidation numbers, although some are positive in certain compounds or ions. • Fluorine always has an oxidation number of −1. • The other halogens have an oxidation number of −1 when they are negative; they can have positive oxidation numbers, however, most notably in oxyanions.

  19. Assigning Oxidation Numbers • The sum of the oxidation numbers in a neutral compound is 0. • The sum of the oxidation numbers in a polyatomic ion is the charge on the ion.

  20. What is the oxidation state of manganese in the permanganate ion? MnO4-1 • +8 • -1 • +2 • +7

  21. What is the oxidation state of sulfur in the H2SO4? • +6 • +4 • +8 • -6

  22. What is the oxidation state of chlorine in HCl? • +1 • 0 • -1 • Can’t tell

  23. What is the oxidation state of chlorine in the HClO4? • 0 • +4 • +7 • -7

  24. What chemical process is happening to manganese in the following reaction? MnO4− + C2O42- Mn2+ + CO2 • oxidation • reduction • Both 1 and 2 • None of the above

  25. What chemical process is happening to oxygen in the following reaction? MnO4− + C2O42- Mn2+ + CO2 • oxidation • reduction • Both 1 and 2 • None of the above

  26. What chemical process is happening to carbon in the following reaction? MnO4− + C2O42- Mn2+ + CO2 • oxidation • reduction • Both 1 and 2 • None of the above

  27. Balancing Oxidation-Reduction Equations Perhaps the easiest way to balance the equation of an oxidation-reduction reaction is via the half-reaction method.

  28. Balancing Oxidation-Reduction Equations This involves treating (on paper only) the oxidation and reduction as two separate processes, balancing these half reactions, and then combining them to attain the balanced equation for the overall reaction.

  29. Half-Reaction Method • Assign oxidation numbers to determine what is oxidized and what is reduced. • Write the oxidation and reduction half-reactions.

  30. Half-Reaction Method • Balance each half-reaction. • Balance elements other than H and O. • Balance O by adding H2O. • Balance H by adding H+. • Balance charge by adding electrons. • Multiply the half-reactions by integers so that the electrons gained and lost are the same.

  31. Half-Reaction Method • Add the half-reactions, subtracting things that appear on both sides. • Make sure the equation is balanced according to mass. • Make sure the equation is balanced according to charge.

  32. Half-Reaction Method Consider the reaction between MnO4− and C2O42− : MnO4−(aq) + C2O42−(aq) Mn2+(aq) + CO2(aq)

  33. +7 +3 +2 +4 MnO4− + C2O42- Mn2+ + CO2 Half-Reaction Method First, we assign oxidation numbers. Since the manganese goes from +7 to +2, it is reduced. Since the carbon goes from +3 to +4, it is oxidized.

  34. Oxidation Half-Reaction C2O42− CO2 To balance the carbon, we add a coefficient of 2: C2O42− 2 CO2

  35. Oxidation Half-Reaction C2O42− 2 CO2 The oxygen is now balanced as well. To balance the charge, we must add 2 electrons to the right side. C2O42− 2 CO2+ 2 e−

  36. Reduction Half-Reaction MnO4− Mn2+ The manganese is balanced; to balance the oxygen, we must add 4 waters to the right side. MnO4− Mn2++ 4 H2O

  37. Reduction Half-Reaction MnO4− Mn2+ + 4 H2O To balance the hydrogen, we add 8 H+ to the left side. 8 H+ + MnO4− Mn2+ + 4 H2O

  38. Reduction Half-Reaction 8 H+ + MnO4− Mn2+ + 4 H2O To balance the charge, we add 5 e− to the left side. 5 e− + 8 H+ + MnO4− Mn2+ + 4 H2O

  39. Combining the Half-Reactions Now we evaluate the two half-reactions together: C2O42− 2 CO2 + 2 e− 5 e− + 8 H+ + MnO4− Mn2+ + 4 H2O To attain the same number of electrons on each side, we will multiply the first reaction by 5 and the second by 2.

  40. Combining the Half-Reactions 5 C2O42− 10 CO2 + 10 e− 10 e− + 16 H+ + 2 MnO4− 2 Mn2+ + 8 H2O When we add these together, we get: 10 e− + 16 H+ + 2 MnO4− + 5 C2O42−  2 Mn2+ + 8 H2O + 10 CO2 +10 e−

  41. Combining the Half-Reactions 10 e− + 16 H+ + 2 MnO4− + 5 C2O42−  2 Mn2+ + 8 H2O + 10 CO2 +10 e− The only thing that appears on both sides are the electrons. Subtracting them, we are left with: 16 H+ + 2 MnO4− + 5 C2O42−  2 Mn2+ + 8 H2O + 10 CO2

  42. Balancing in Basic Solution • If a reaction occurs in basic solution, one can balance it as if it occurred in acid. • Once the equation is balanced, add OH− to each side to “neutralize” the H+ in the equation and create water in its place. • If this produces water on both sides, you might have to subtract water from each side.

  43. Activity series • Some elements are more reactive than others. • What factors make elements more reactive? • Atomic radius (shielding/nuclear charge) • Stability • Activity series shows the order of reactivity

  44. Do metals… • Lose electrons? • Gain electrons? • Both of the above • Neither 1 nor 2

  45. Which alkali metal is the most reactive? • Lithium • Sodium • Potassium • Rubidium

  46. Activity Series • Activity series is usually written in terms of species most likely to be oxidized • Elements higher on the chart want to be oxidized more than elements lower on the chart. • Example: Potassium wants to be oxidized more than Sodium • What elements want to be oxidized more than Silver? • In Zumdahl, it is similar to the Standard Reduction Potential chart which is written in terms of reduction • Species more likely to be reduced are less likely to be oxidized • So in Zumdahl, the chart is in the opposite order

  47. Activity Series • Activity series allow you to predict whether a redox reaction will happen or not. • Elements higher on the chart will displace a metal ion of an element lower on the chart. • Mg + Zn2+ will react to form Zn and Mg2+

  48. Will K reduce Na+? • Yes • No • Need more information

  49. Will Mg2+ reduce Na+? • Yes • No • Need more information

  50. Will Mg reduce Na+? • Yes • No • Need more information

More Related