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, O 2, Cl 2, S 2. The oxidation number of a monatomic ion is equal to the charge on the ion. – Ex: the oxidation number of a Ca 2+ 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 SiCl 4, 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 H 2 O 2, 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, CaH 2, and AlH 3 ; 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 KClO 3 (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) = 0X = +5 What is the oxidation number of sulfur in SO 3 2– (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) = -2X = +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. Another Way to Remember OIL RIG xidationxidationxidationxidation s oseoseoseose eductioneductioneductioneduction s ainainainain

15. Oxidation and Reduction Zinc is oxidized from zinc metal to the Zn 2+ ion. – H + is the oxidizing agent. Each H + is reduced and combine to form H 2. – 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 (H 2 O 2 ) 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 1)Assign oxidation numbers to determine what is oxidized and what is reduced. 2)Write the oxidation and reduction half-reactions. 3)Balance each half-reaction. – Balance elements other than H and O. – Balance O by adding H 2 O. – Balance H by adding H+. – Balance charge by adding electrons. 4)Multiply the half-reactions by integers so that the electrons gained and lost are the same.

24. Half-Reaction Method 5)Add the half-reactions, subtracting things that appear on both sides. 6)Make sure the equation is balanced according to mass. 7)Make sure the equation is balanced according to charge.

25. Half-Reaction Method Consider the reaction between MnO 4− and C 2 O 4 2− : –MnO 4 − (aq) + C 2 O 4 2− (aq)  Mn 2+ (aq) + CO 2 (g)

26. Half-Reaction Method First, we assign oxidation numbers. MnO 4 − + C 2 O 4 2-  Mn 2+ + CO 2 +7+3+4+2 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 C 2 O 4 2−  CO 2 To balance the carbon, we add a coefficient of 2: C 2 O 4 2−  2CO 2 The oxygen is now balanced as well. To balance the charge, we must add 2 electrons to the right side. C 2 O 4 2−  2CO 2 + 2e−

28. Reduction Half-Reaction MnO 4 −  Mn 2+ The manganese is balanced; to balance the oxygen, we must add 4 waters to the right side. MnO 4 −  Mn 2+ + 4H 2 O To balance the hydrogen, we add 8 H+ to the left side. 8H + + MnO 4 −  Mn 2+ + 4H 2 O To balance the charge, we add 5 e− to the left side. 5e − + 8H + + MnO 4 −  Mn 2+ + 4H 2 O

29. Combining the Half-Reactions Now we evaluate the two half-reactions together: C 2 O 4 2−  2CO 2 + 2e − 5e − + 8H + + MnO 4 −  Mn 2+ + 4H 2 O 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 5C 2 O 4 2−  10CO 2 + 10e − 10e − + 16H + + 2MnO 4 −  2Mn 2+ + 8H 2 O When we add these together, we get: 10e − + 16H + + 2MnO 4 − + 5C 2 O 4 2−  2Mn 2+ + 8H 2 O + 10CO 2 +10e − The only thing that appears on both sides are the electrons. Subtracting them, we are left with: 16H + + 2MnO 4 − + 5C 2 O 4 2−  2Mn 2+ + 8H 2 O + 10CO 2

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 (E 0 ) 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 E 0 the greater the tendency for the substance to be reduced The half-cell reactions are reversible – The sign of E 0 changes when the reaction is reversed Changing the stoichiometric coefficients of a half-cell reaction does not change the value of E 0

39. Standard Electrode Potentials Standard emf (E 0 cell) E 0 = E 0 cathode + E 0 anode If the reaction is backwards, be sure to flip the sign!

40. Cell Diagram Zn (s) + Cu 2+ (aq) Cu (s) + Zn 2+ (aq) [Cu 2+ ] = 1 M & [Zn 2+ ] = 1 M Zn (s) | Zn 2+ (1 M) || Cu 2+ (1 M) | Cu (s) anodecathode Determining Cell Potential The difference in electrical potential between the anode and cathode is called: – cell voltage – electromotive force (emf) – cell potential

41. Why Study Electrochemistry? Batteries Corrosion Industrial production of chemicals such as Cl 2, NaOH, F 2 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  Zn 2+ + 2e- Cathode (+) – 2NH 4 + + 2e-  2NH 3 + H 2

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 reactions involve 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 Reactions involve electrons, not protons and neutrons

50. 23.1 Comparison

51. Radioactivity It is not uncommon for some nuclides of an element to be unstable, or radioactive. We refer to these as radionuclides. – There are several ways radionuclides can decay into a different nuclide.

52. Types of Radiation An alpha particle (α) has the same composition as a helium nucleus—two protons and two neutron. – The charge of an alpha particle is 2+ due to the presence of the two protons. – Alpha radiation consists of a stream of alpha particles.

53. Alpha Decay All nuclei with more than 83 protons are radioactive and decay spontaneously. These very heavy nuclei often decay by emitting alpha particles.

54. Types of Radiation Because of their mass and charge, alpha particles are relatively slow- moving compared with other types of radiation. A beta particle is a very-fast moving electron that has been emitted from a neutron of an unstable nucleus.

55. Types of Radiation Note that the mass number of the product nucleus is the same as that of the original nucleus (they are both 131), but its atomic number has increased by 1 (54 instead of 53). – This change in atomic number, and thus, change in identity, occurs because the electron emitted during the beta decay has been removed from a neutron, leaving behind a proton.

56. Beta Decay Beta Particle Production – β particle is an electron Can assume the mass is zero – Net effect : changing a n into a p How can an electron come from a nucleus? Loss of a  -particle (a high energy electron)  0 −1 e 0 −1 or I 131 53 Xe 131 54  + e 0 −1

57. Gamma Emission Gamma Ray Production – γ ray is collection of high energy photons – Occurs with other types of decay – Helps a nucleus release extra energy so it can relax to a lower energy state  0000

58. Penetrating Ability

59. Positron Emission Positron Production – Occurs for nuclides below the line of stability – Positron is a positive particle with same mass of electron Also called antiparticle of electron Net effect: change p into n e 0101 C 11 6  B 11 5 + e 0101

60. Electron Capture (K-Capture) Electron Capture – One of the inner electrons in an atom is captured by a proton in the nucleus As a result, a proton is transformed into a neutron. – Gamma rays always produced Decay Series – When several types of decay occur until a stable nuclide is produced p 1111 + e 0 −1  n 1010

61. Geiger Counter Used to detect radioactive substances

62. Nuclear Stability Nuclear force is a force that acts on subatomic particles overcoming the repulsion between protons. – For atoms with low atomic numbers (< 20), stable nuclei have neutron-to- proton ratios of 1 : 1. – As atomic number increases, more and more neutrons are needed to balance the electrostatic repulsion forces.

63. Radioactivity and Stability All nuclides are unstable with 84 p or more – Lightweight nuclides are stable with equal numbers of neutrons and protons – Heavy nuclides should have ratio >1 to be stable

64. Stable Numbers Certain numbers of neutrons and protons are extra stable – n or p = 2, 8, 20, 50, 82 and 126 – Like extra stable numbers of electrons in noble gases (e- = 2, 10, 18, 36, 54 and 86) Nuclei with even numbers of both protons and neutrons are more stable than those with odd numbers of neutron and protons All isotopes of the elements with atomic numbers higher than 83 are radioactive All isotopes of Tc and Pm are radioactive

65. Neutron - Proton Ratios Any element with more than one proton (i.e., anything but hydrogen) will have repulsions between the protons in the nucleus. A strong nuclear force helps keep the nucleus from flying apart. Neutrons play a key role stabilizing the nucleus. – Therefore, the ratio of neutrons to protons is an important factor. For smaller nuclei (Z  20) stable nuclei have a neutron-to-proton ratio close to 1:1. As nuclei get larger, it takes a greater number of neutrons to stabilize the nucleus.

66. Stable Nuclei The shaded region in the figure shows what nuclides would be stable, the so-called belt of stability. Nuclei above this belt have too many neutrons. – They tend to decay by emitting beta particles. Nuclei below the belt have too many protons. – They tend to become more stable by positron emission or electron capture. There are no stable nuclei with an atomic number greater than 83. – These nuclei tend to decay by alpha emission.

67. Radioactive Series Large radioactive nuclei cannot stabilize by undergoing only one nuclear transformation. – They undergo a series of decays until they form a stable nuclide (often a nuclide of lead).

68. Nuclear Transformations Nuclear transformations are a Change of one element into another – Scientists have been able to use this to make the periodic table larger by creating new elements Since 1940, have been able to make transuranium elements (93-112)

69. Half-Life Half-Life is the time that it takes for 1/2 a sample to decompose. – The rate of a nuclear transformation depends only on the “reactant” concentration. Decay of 20.0 mg of 15 O. What remains after 3 half-lives? After 5 half-lives?

70. Radioactive Decay Rates In the equation, n is equal to the number of half- lives that have passed.

71. Nuclear Fission and Fusion Fission is the s plitting a heavy nucleus into 2 smaller nuclei with smaller mass numbers – Can use neutrons to create instability – Neutrons produced are used to cause more fission – Produces a huge amount of energy

72. Fission NEUTRON U-235 NUCLEUS Kr-92 NUCLEUS Ba-141 NUCLEUS NEUTRONS

73. Nuclear Fission How does one tap all that energy? Nuclear fission is the type of reaction carried out in nuclear reactors. – Bombardment of the radioactive nuclide with a neutron starts the process. – Neutrons released strike other nuclei, causing their decay and the production of more neutrons. This process continues in what we call a nuclear chain reaction.

74. Nuclear Fission If there are not enough radioactive nuclides in the path of the ejected neutrons, the chain reaction will die out. – Therefore, there must be a certain minimum amount of fissionable material present for the chain reaction to be sustained: Critical Mass.

75. Nuclear Reactors In nuclear reactors the heat generated by the reaction is used to produce steam that turns a turbine connected to a generator.

76. Nuclear Reactors The reaction is kept in check by the use of control rods. – These block the paths of some neutrons, keeping the system from reaching a dangerous supercritical mass.

77. Fusion + + H H H H 1111 1111 1111 He 2121 2121 4242 1111 H H

78. Nuclear Fusion Fusion – Combination of 2 light nuclei to form a heavier, more stable nucleus – Stars produce their energy using this – Requires very high temperatures – Must be shot at each other to get close enough

79. Nuclear Fusion Fusion would be a superior method of generating power. – The good news is that the products of the reaction are not radioactive. – The bad news is that in order to achieve fusion, the material must be in the plasma state at several million kelvins. Tokamak apparati like the one shown at the right show promise for carrying out these reactions.

80. 2H + 3H 4He + 1n + 1 1 2 0 Occurs in the sun and other stars Nuclear Fusion Energy Fusion –Excessive heat can not be contained –Attempts at “cold” fusion have FAILED. –“Hot” fusion is difficult to contain

81. Effects of Radiation Any sort of energy transferred to cells can break bonds and cause damage Radioactive species are sources of high energy particles so can be very harmful Types – Somatic: cause illness, cancer, death – Genetic: produce damage in offspring

82. Factors in Effects of Radiation The more energy, the more damage – How deep it goes into body γ rays > β particles (1 cm) > α particles (skin) – How easily they attract electrons from biomolecules (ionization) γ rays cause less than α particles – How long it stays inside body

83. Medical Applications Radiotracers – Radioactive nuclides that can be traced in people by monitoring their radioactivity Thallium-201 – For assessing heart damage from heart attacks – Is taken up by healthy heart tissue only

84. Medical Applications Iodine-131 – For diagnosing thyroid problems – Patients drink a solution of 131-I and the uptake is monitored

85. End of Unit 11. Be Prepared for Unit 11 Test.