Unit 11: RedOx

1. Unit 11: RedOx Chapter 20 & 21 Pre-AP Chemistry I Edmond North High School

2. Electron Transfer Reactions • Electron transfer reactions are oxidation-reduction or redox reactions. • Results in the generation of an electric current (electricity) or be caused by imposing an electric current. • Therefore, this field of chemistry is often called ELECTROCHEMISTRY.

3. What’s the Point of RedOx? • REDOX reactions are important in … • Biological Processes • Electrical production (batteries, fuel cells) • Electroplating metals • Protecting metals from corrosion • Balancing complex chemical equations • Sensors and machines (e.g. pH meter)

4. Oxidation Number • The oxidation number of an atom is the number of electrons lost or gained when it forms ions. • Oxidation numbers are written with the sign before the number, whereas ionic charge is written after the number. • Oxidation number: +3 • Ionic charge: 3+

5. Rules for Oxidation Numbers 1. The oxidation number of an uncombined atom is zero. • Ex: Mg, Ca, O2, Cl2, S 2. The oxidation number of a monatomic ion is equal to the charge on the ion. • Ex: the oxidation number of a Ca2+ is +2, and Br– is –1.

6. Rules for Oxidation Numbers 3. The oxidation number of the more electronegative atom in a molecule or a polyatomic ion is the charge of its ion. • In SiCl4, chlorine is more electronegative, so chlorine has an oxidation number of –1. 4. The most electronegative element, fluorine, always has an oxidation number of –1 when it is bonded to another element.

7. Rules for Oxidation Numbers 5. The oxidation number of oxygen in compounds is –2 • Exceptions: • Peroxides, such as H2O2, it is –1. • When bonded to fluorine, the oxidation number is +2 6. The oxidation number of hydrogen in most of its compounds is +1. • Exception: when hydrogen bonds as an anion such as LiH, CaH2, and AlH3; its oxidation number is –1.

8. Rules for Oxidation Numbers 7. The sum of the oxidation numbers in a neutral compound is zero. 8. The sum of the oxidation numbers of the atoms in a polyatomic ion is equal to the charge on the ion.

9. Common Oxidation Numbers

10. Determining Oxidation Numbers Practice • What is the oxidation number of chlorine in KClO3 (potassium chlorate) • Neutral salt, so oxidation numbers must add up to zero. • Rule 5, the oxidation number of oxygen in compounds is –2. • Rule 7 states Group 1 metals have a +1 oxidation number. • (+1) + x + 3(-2) = 0 X = +5 • What is the oxidation number of sulfur in SO32– (sulfite ion) • Ion has a charge of 2–, so oxidation numbers must add up to –2. • Rule 5, the oxidation number of oxygen in compounds is –2. • X + 3(-2) = -2 X = +4

11. Redox Reactions • RedOx (oxidation-reduction) reactions occur when oxidation numbers change.

12. Terminology for Redox • OXIDATION - loss of electron(s) by a species; increase in oxidation number; increase in oxygen. • REDUCTION - gain of electron(s); decrease in oxidation number; decrease in oxygen; increase in hydrogen. • OXIDIZING AGENT - electron acceptor; species is reduced. (an agent facilitates something; ex. Travel agents don’t travel, they facilitate travel) • REDUCING AGENT - electron donor; species is oxidized.

13. You Can’t Have One Without the Other! • Reduction (gaining electrons) can’t happen without an oxidation to provide the electrons. • You can’t have 2 oxidations or 2 reductions in the same equation. Reduction has to occur at the cost of oxidation LEO the lion says GER! GER!

14. OIL RIG Another Way to Remember s s xidation ose eduction ain

15. Oxidation and Reduction • Zinc is oxidized from zinc metal to the Zn2+ ion. • H+ is the oxidizing agent. • Each H+ is reduced and combine to form H2. • Zn is the reducing agent.

16. Oxidation Number in Redox Reactions • To see how oxidation numbers change, start by assigning numbers to all elements in the balanced equation. • There is no change in the oxidation number of potassium. • The potassium ion takes no part in the reaction and is called a spectator ion.

17. Oxidizing and Reducing Agents • Oxidizing and reducing agents play significant roles in your daily life. • For example, when you add bleach to your laundry, you are using sodium hypochlorite (NaClO), an oxidizing agent. • Hydrogen peroxide (H2O2) can be used as an antiseptic because it oxidizes some of the vital biomolecules of germs.

18. Oxidation–Reduction Reactions Practice • Identify what is oxidized and what is reduced in this reaction. • Aluminum is oxidized, Iron is reduced • Identify the oxidizing agent and the reducing agent. • Aluminum is the reducing agent, Iron is the oxidizing agent.

19. Equations Must Balanced • There are two conditions now for molecular, ionic, and net ionic equations • Mass Balance • Both sides of an equation should have the same number of each type of atom • Charge Balance • Both sides of a reaction should have the same net charge

20. Half-Reactions • The oxidation process and the reduction process of a redox reaction can each be expressed as a half-reaction. • For example, consider the unbalanced equation for the formation of aluminum bromide. • This is a method for tracking RedOx on PAPER ONLY!

21. Half-Reactions • The oxidation half-reaction shows the loss of electrons by aluminum. • The reduction half-reaction shows the gain of electrons by bromine.

22. Balancing RedOx Equations • Perhaps the easiest way to balance the equation of an oxidation-reduction reaction is via the half-reaction method. • 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.

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

24. 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.

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

26. +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.

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

28. Reduction Half-Reaction MnO4− Mn2+ • The manganese is balanced; to balance the oxygen, we must add 4 waters to the right side. MnO4− Mn2+ + 4H2O • To balance the hydrogen, we add 8 H+ to the left side. 8H+ + MnO4− Mn2+ + 4H2O • To balance the charge, we add 5 e− to the left side. 5e−+ 8H+ + MnO4− Mn2+ + 4H2O

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

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

31. Spontaneous RedOx • In spontaneous oxidation-reduction (redox) reactions, electrons are transferred and energy is released.

32. Electric Current • To obtain an electrical current, we separate the oxidizing and reducing agents so electron transfer occurs thru an external wire. • This is accomplished in an electrochemical cell. • A group of such cells is called a battery.

33. Electrochemical Cells • A typical cell looks like this. • The oxidation occurs at the anode. • The reduction occurs at the cathode. • Refer to your activity series to determine anode and cathode. • Most active metal is oxidized (anode).

34. Activity Series

35. Electrochemical Cells • Electrons leave the anode and flow through the wire to the cathode. • As the electrons leave the anode, the cations formed dissolve into the solution in the anode compartment. • Once even one electron flows, the charges in each beaker would not be balanced and the flow of electrons would stop.

36. Electrochemical Cells • We use a salt bridge, usually a U-shaped tube that contains a salt solution, to keep the charges balanced. • Cations move toward the cathode. • Anions move toward the anode.

37. Standard Electrode Potentials • Standard reduction potential (E0) is the voltage associated with a reduction reaction at an electrode when all solutes are 1 M and all gases are at 1 atm.

38. Standard Reduction Potentials • The more positive E0 the greater the tendency for the substance to be reduced • The half-cell reactions are reversible • The sign of E0 changes when the reaction is reversed • Changing the stoichiometric coefficients of a half-cell reaction does not change the value of E0

39. Standard Electrode Potentials • Standard emf (E0 cell) • E0 = E0cathode + E0anode • If the reaction is backwards, be sure to flip the sign!

40. Determining Cell Potential • The difference in electrical potential between the anode and cathode is called: • cell voltage • electromotive force (emf) • cell potential Cell Diagram Zn (s) + Cu2+(aq) Cu (s) + Zn2+(aq) [Cu2+] = 1 M & [Zn2+] = 1 M Zn (s) | Zn2+ (1 M) || Cu2+ (1 M) | Cu (s) anode cathode

41. Why Study Electrochemistry? • Batteries • Corrosion • Industrial production of chemicals such as Cl2, NaOH, F2 and Al • Biological redox reactions The heme group

42. Photosynthesis is RedOx • The process that uses the sun’s energy to transfer electrons to make glucose during photosynthesis.

43. Cellular Respiration is RedOx • The process by which food molecules breakdown to produce ATP is called cellular respiration. • The last two stage is the electron transport chain and is a series of redox reactions with oxygen as the final oxidizing agent.

44. Dry Cell Battery • Anode (-) • Zn  Zn2+ + 2e- • Cathode (+) • 2NH4+ + 2e-  2NH3 + H2

45. Statue of Liberty • Why is the Statue of Liberty green • Oxidation of Copper! • As copper oxidizes it turns to copper oxide which has a green color.

46. The Titanic • A rusticle is a formation of rust similar to an icicle or stalactite in appearance that occurs underwater when iron oxidizes. • They may be familiar from underwater photographs of shipwrecks.

47. Electrolysis • Electrolysis is running a galvanic cell backwards. • Put a voltage bigger than the potential and reverse the direction of the redox reaction. • Used for electroplating.

48. Radioactivity One of the pieces of evidence for the fact that atoms are made of smaller particles came from the work of Marie Curie(1876-1934). She discovered radioactivity, the spontaneous disintegration of some elements into smaller pieces.

49. Nuclear Reactions vs. Normal Chemical Changes Nuclear reactionsinvolve the nucleus The nucleus opens, and protons and neutrons are rearranged The opening of the nucleus releases a tremendous amount of energy that holds the nucleus together – called binding energy “Normal” Chemical Reactionsinvolve electrons, not protons and neutrons

50. Comparison 23.1