Unit 11: RedOx and Nuclear Chemistry

1. Unit 11: RedOx and Nuclear Chemistry Chapters 20 & 21 General Chemistry 1 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 for 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. Spontaneous RedOx • In spontaneous oxidation-reduction (redox) reactions, electrons are transferred and energy is released.

23. 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.

24. 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).

25. Activity Series

26. 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.

27. 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.

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

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

30. 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.

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

32. 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.

33. 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.

34. Nuclear Radiation Nuclear chemistryis the study of the structure of atomic nuclei and the changes they undergo. In 1895, Wilhelm Roentgen(1845–1923) found that invisible rays were emitted when electrons bombarded certain materials. The emitted rays were discovered because they caused photographic plates to darken. Roentgen named these invisible high-energy emissions X rays.

35. The Discovery of Radioactivity Marie Curie(1867–1934) and her husband Pierre (1859–1906) took mineral samples (called pitchblende) and isolated the components emitting the rays. Marie Curie named the process by which materials give off such raysradioactivity; the rays and particles emitted by a radioactive source are called radiation.

36. Types of Radiation Unstable nuclei emit radiation to attain more stable configurations in a process called radioactive decay. The three most common types of radiation are alpha (α), beta (β), and gamma (γ).

37. 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 radiationconsists of a stream of alpha particles.

38. Balancing a Nuclear Equation Write a balanced nuclear equation for the alpha decay of thorium-230. Thorium-230 is the initial reactant, while the alpha particle is one of the products of the reaction. The reaction is summarized below. You must determine the unknown product of the reaction, X. This can be done through the conservation of atomic number and mass number.

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

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

41. 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.

42. Types of Radiation Gamma raysare high-energy (short wavelength) electromagnetic radiation. Gamma rays almost always accompany alpha and beta radiation. Because gamma rays have no effect on mass number or atomic number, it is customary to omit them from nuclear equations.

43. Penetrating Ability

44. Nuclear Stability Nuclear forceis 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.

45. Radioactive Decay Rates Radioactive decay rates are measured in half-lives. A half-life is the time required for one-half of a radioisotope’s nuclei to decay into its products. For example, the half-life of the radioisotope strontium-90 is 29 years. If you had 10.0 g of strontium-90 today, 29 years from now you would have 5.0 g left.

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

47. Radiochemical Dating • The process of determining the age of an object by measuring the amount of a certain radioisotope remaining in that object is called radiochemical dating.

48. Nuclear Fission The splitting of a nucleus into fragments is called nuclear fission. Nuclear fission releases a large amount of energy. One fission reaction can lead to more fission reactions, a process called a chain reaction. The mass needed to sustain a chain reaction is called critical mass.

49. Nuclear Reactors Nuclear power plantsuse the process of nuclear fission to produce heat in nuclear reactors. The heat generates steam, which drives turbines that produce electricity. Fissionable uranium (IV) oxide (UO2) is used as fuel. Cadmium and boronare used to keep the fission process under control.

50. Nuclear Fusion The combining of atomic nuclei is called nuclear fusion. For example, nuclear fusion occurs within the Sun, where hydrogen atoms fuse to form helium atoms. Fusion reactionscan release very large amounts of energy but require extremely high temperatures. For this reason, they are also called thermonuclear reactions.