1. PACKET #12: Oxidation/Reduction (REDOX) Reference Table: Periodic Table & J www.regentsprep.org

2. Before we start, recall... each box on the periodic table has one or multiple oxidation numbers on the upper right hand corner of the box. each box on the periodic table has one or multiple oxidation numbers on the upper right hand corner of the box. The oxidation number is the charge that an atom takes on after either losing or gaining one or more electrons in order to most efficiently stabilize its outer most shell (act like a Noble Gas) using the least amount of energy. The oxidation number is the charge that an atom takes on after either losing or gaining one or more electrons in order to most efficiently stabilize its outer most shell (act like a Noble Gas) using the least amount of energy. Any free element (not combined with any other element) has an oxidation number of 0. Any free element (not combined with any other element) has an oxidation number of 0. Ex: Na = Na0

3. Oxidation loss of electrons by a molecule, atom, or ion loss of electrons by a molecule, atom, or ion Remember you learned that metals lose electrons to have a complete outer shell (more stable) Remember you learned that metals lose electrons to have a complete outer shell (more stable) For example Na will lose one electron to become Na+ ion For example Na will lose one electron to become Na+ ion Therefore Na is Oxidized Therefore Na is Oxidized When any atom loses electrons its oxidation number INCREASES When any atom loses electrons its oxidation number INCREASES Example: Na 0  Na + + e - -its oxidation # increases from 0 to + 1

4. Reduction The gain of electrons by a molecule, atom, or ion The gain of electrons by a molecule, atom, or ion Nonmetals gain electron to have a complete outer shell Nonmetals gain electron to have a complete outer shell Phosphorous will gain three electrons to have 8 in its outer shell Phosphorous will gain three electrons to have 8 in its outer shell Because P gains 3 electrons P is Reduced Because P gains 3 electrons P is Reduced When any atoms gains electrons (negative particles) the oxidation number DECREASES When any atoms gains electrons (negative particles) the oxidation number DECREASES Example: P 0 + 3e -  P -3 -its oxidation # decreases from 0 to - 3

5. JUST REMEMBER... “OIL RIG” (oxidation is loss, reduction is gain)

6. REDOX Short-hand for an oxidation/reduction equation. In a single reaction there is both oxidation and reduction. Short-hand for an oxidation/reduction equation. In a single reaction there is both oxidation and reduction. REDOX reactions have conservation of matter, and conservation of charge. The numbers of elements are equal on both the reactant and the product side, and the total charge on both the reactant and product side equal 0. REDOX reactions have conservation of matter, and conservation of charge. The numbers of elements are equal on both the reactant and the product side, and the total charge on both the reactant and product side equal 0. Na 0  Na + + e - P 0 + 3e -  P -3 3Na + P  N 3 P (sodium phosphide) Na is being oxidized; P is being reduced Electrons to the left of the arrow means that they are being gained, to the right of the arrow means they are being lost.

7. REDOX In a REDOX reaction, when something is being oxidized, it is called the reducing agent, and when something is being reduced it is called the oxidizing agent In a REDOX reaction, when something is being oxidized, it is called the reducing agent, and when something is being reduced it is called the oxidizing agent Reducing Agent: an electron donor Reducing Agent: an electron donor Oxidizing Agent: an electron acceptor Oxidizing Agent: an electron acceptor

8. Oxidation States Oxidation States are assigned to atoms to identify how many electrons are either gained or lost by an atom Oxidation States are assigned to atoms to identify how many electrons are either gained or lost by an atom For example metals in Group 2 (like Ca) have a +2 oxidation state. Therefore Metals in Group 2 lose two electrons when they form compounds For example metals in Group 2 (like Ca) have a +2 oxidation state. Therefore Metals in Group 2 lose two electrons when they form compounds Changes in oxidation numbers indicate that a redox reaction has occurred. Changes in oxidation numbers indicate that a redox reaction has occurred. It is important to learn the rules for assigning oxidation states to atoms in order to determine whether oxidation or reduction has occurred. It is important to learn the rules for assigning oxidation states to atoms in order to determine whether oxidation or reduction has occurred.

9. Oxidation State Rules 1) Free elements (not combined with any other element) have an oxidation number of zero. Ex: Na, O 2, H 2 2) All metals in Group 1 have an oxidation number of +1. 3) All metals in Group 2 have an oxidation number of +2. 4) F (fluorine) always has an oxidation of -1

10. Oxidation State Rules 5) The oxidation of simple ions is equal to the charge on the ion. Ex: Mg+2 has an oxidation number of +2. 6) The sum of the oxidation numbers must equal 0 Examples: NaCl & MgCl 2 7) In polyatomic ions, the sum of the oxidation numbers of all the atoms must equal the charge of the ion. Example: sulfate ion SO 4 -2. Oxygen has an oxidation of -2, and therefore (-2) x (+4) = -8. Remember that the overall charge of this ion has to be -2. What is the oxidation # of sulfur?

11. Oxidation State Rules 8) In general, oxygen has an oxidation number of -2. Oxygen has an oxidation number of -1 in peroxides (O 2 -2 ) Example: H 2 O 2. Oxygen has an oxidation number of +2 in compounds with fluorine Example: OF 2 9) Hydrogen has an oxidation number of +1 in all compounds combined with a non-metal. Hydrogen has an oxidation number of -1 when it is a metal hydrides (metal and hydrogen. Example: LiH, and CaH 2

12. Assigning Oxidation Numbers When you have a binary compound or polyatomic compound, assigning oxidation numbers is a little easier than when you have to assign oxidation numbers to compounds with more than two elements. When you have a binary compound or polyatomic compound, assigning oxidation numbers is a little easier than when you have to assign oxidation numbers to compounds with more than two elements. EXAMPLES: NaClCaO NaClCaO CO 2 Li 3 N CO 2 Li 3 N

13. Assigning Oxidation Numbers Dr. McGuiness’ “bookend” technique to assigning oxidation numbers to compounds with more than two elements. Dr. McGuiness’ “bookend” technique to assigning oxidation numbers to compounds with more than two elements. 1. Identify the oxidation # of the last element (overall charge) 2. Identify the oxidation # of the first element (overall charge) 3. If there is no charge to the compound, then the overall charge must be 0, therefore you can determine the oxidation # of the element in the middle. Li(MnO 4 ) (+1) + (?) + (-8) = 0 -2 x 4 = -8 +1 x 1 = +1

15. Half-Reactions Once you can assign oxidation numbers, then you can take a REDOX reaction and break it down to half-reactions in order to figure out what is being oxidized and what is being reduced. Once you can assign oxidation numbers, then you can take a REDOX reaction and break it down to half-reactions in order to figure out what is being oxidized and what is being reduced. Half-reaction: shows either oxidation or reduction. A redox reaction is made up of two half reactions (one oxidation, and one reduction). Half-reaction: shows either oxidation or reduction. A redox reaction is made up of two half reactions (one oxidation, and one reduction). A redox reaction is when the oxidation number of the half-reaction and the reduction half-reaction occur simultaneously. A redox reaction is when the oxidation number of the half-reaction and the reduction half-reaction occur simultaneously.

16. Half-Reactions Example: Cu + Ag(NO 3 ) → Cu(NO 3 ) 2 + Ag 1. Assign oxidation numbers to everything 2. See which oxidation numbers change from reactant side to product side. 3. Determine which half-reaction is oxidation (loses e - ), and which is reduction (gains e - ) OX: Cu 0 → Cu 2+ + 2e - OX: Cu 0 → Cu 2+ + 2e - RED: Ag + + 1e - → Ag 0 RED: Ag + + 1e - → Ag 0

17. Balancing REDOX reactions Write half reactions for oxidation and reduction. Write half reactions for oxidation and reduction. Balance the half reactions separately. Balance the half reactions separately. Equalize the number of electrons gained and lost. Equalize the number of electrons gained and lost. Add the half reactions together and cancel the electrons to get the final balanced equation. Add the half reactions together and cancel the electrons to get the final balanced equation. The electrons lost in a REDOX reaction have to equal the electrons gained The electrons lost in a REDOX reaction have to equal the electrons gained

18. Balancing REDOX reactions Cu + 2Ag(NO 3 ) → Cu(NO 3 ) 2 + 2Ag OX: Cu 0 → Cu 2+ + 2e - OX: Cu 0 → Cu 2+ + 2e - RED: 2(Ag + + 1e - → Ag)  2Ag + + 2e - 2Ag 0 RED: 2(Ag + + 1e - → Ag)  2Ag + + 2e - 2Ag 0 When copper is oxidized, it loses 2 electrons, which are gained by the silver ion. When copper is oxidized, it loses 2 electrons, which are gained by the silver ion. Copper = reducing agent Copper = reducing agent Silver = oxidizing agent Silver = oxidizing agent

19. Let’s Practice Let’s Practice For the following REDOX reactions complete the following: For the following REDOX reactions complete the following: 1. Write the correct half-reactions 2. Identify what is oxidized/what is reduced 3. Identify the oxidizing/reducing agents 4. Balance the equation

20. 1. Ce 4+ + Sn 2+ → Ce 3+ + Sn 4+ 2. Fe + SnBr 4 → FeBr 2 + Sn 3. Cu + Ag(NO 3 ) → Cu(NO 3 ) 2 + Ag

21. Activity Series – Table J Table J compares how active each metal and nonmetal is. Table J compares how active each metal and nonmetal is. Metals higher up are more active, and replace metals from below them from compounds (remember single replacement???) Metals higher up are more active, and replace metals from below them from compounds (remember single replacement???) In a single replacement, the free element has to be more reactive than the element in compound in order for the reaction to be spontaneous. In a single replacement, the free element has to be more reactive than the element in compound in order for the reaction to be spontaneous. If it isn’t – the reaction does NOT GO!! If it isn’t – the reaction does NOT GO!!

22. Spontaneous or Not?? Li + KCl  Ca + MgCO 3  K + LiCl  F 2 + 2NaCl  Cl 2 + 2NaF 

23. Recall that metals lose electrons, and therefore undergo oxidation Recall that metals lose electrons, and therefore undergo oxidation Metals that are more reactive oxidize easier Metals that are more reactive oxidize easier Recall that non-metals gain electrons, and therefore undergo reduction Recall that non-metals gain electrons, and therefore undergo reduction Non-metals that are more reactive reduce easier Non-metals that are more reactive reduce easier

24. Electrochemical Cells There are two types of electrochemical cells, Voltaic & Electrolytic There are two types of electrochemical cells, Voltaic & Electrolytic These cells rely on REDOX reaction in different ways to either generate energy or to separate elements in compound that would normally not exist on their own in nature. These cells rely on REDOX reaction in different ways to either generate energy or to separate elements in compound that would normally not exist on their own in nature.

25. Voltaic Cell A voltaic cell uses REDOX reactions that are spontaneous to produce electricity A voltaic cell uses REDOX reactions that are spontaneous to produce electricity (chemical energy  electrical energy) A battery is an example of a voltaic cell. A battery is an example of a voltaic cell.

26. Voltaic Cell There are two half-cells in a voltaic cell (anode & cathode) There are two half-cells in a voltaic cell (anode & cathode) Each half-cell is a metal strip called an electrode. Each half-cell is a metal strip called an electrode. There is a wire connecting the two electrodes in which electrons travel through. There is a wire connecting the two electrodes in which electrons travel through. Electrons always travel from the Electrons always travel from the anode  cathode

27. Voltaic Cell There is a salt-bridge connecting the two half cells in order to permit ions to flow between the two half-cells. There is a salt-bridge connecting the two half cells in order to permit ions to flow between the two half-cells. Ions travel through the salt bridge that are connecting the two half cells (a complete or closed circuit) Ions travel through the salt bridge that are connecting the two half cells (a complete or closed circuit)

29. Half-Cells Electrodes Salt Bridge Wire

30. Voltaic Cell The anode is the half-cell where oxidation ALWAYS occurs. The anode is the half-cell where oxidation ALWAYS occurs. In a voltaic cell, the anode is negatively charged ( ANODE/OXIDATION – An Ox) In a voltaic cell, the anode is negatively charged ( ANODE/OXIDATION – An Ox) The anode will always be the The anode will always be the half-cell with the metal electrode that is most reactive (Table J) that is most reactive (Table J)

31. Voltaic Cell The cathode is the half-cell where reduction ALWAYS occurs. The cathode is the half-cell where reduction ALWAYS occurs. In a voltaic cell, the cathode is positively charged (CATHODE/REDUCTION - Red Cat) In a voltaic cell, the cathode is positively charged (CATHODE/REDUCTION - Red Cat)

32. Voltaic Cell Let’s consider this REDOX reaction. Let’s consider this REDOX reaction. Zn + Cu +2  Cu + Zn +2 After writing the half-reactions, we can determine which half cell is the anode, and which is the cathode. After writing the half-reactions, we can determine which half cell is the anode, and which is the cathode. Also, Zn is more reactive Also, Zn is more reactive than Cu on Table J.

33. Half-Cell Potential & Cell Voltage Now that we have learned how to “construct” a voltaic cell from Zn and Cu, our next question is “How good does it work?” We measure the output in volt (V). Normally 1.10 V drops to 0 and battery needs to be replaced A list of Standard Electrode Potentials can be used to calculate the maximum standard voltage of a voltaic cell.

35. A standard reduction potential (E°) has been assigned to each reduction half reaction on the list. The more positive - the more easily a the reduction half reaction can occur. A standard reduction potential (E°) has been assigned to each reduction half reaction on the list. The more positive - the more easily a the reduction half reaction can occur. The standard oxidation potential is simply the standard reduction potential with its sign reversed. The standard oxidation potential is simply the standard reduction potential with its sign reversed. Cu +2 + 2e -  Cu(s)E° = +0.34 V Cu(s)  Cu(aq) + 2e - E° = -0.34 V Cu(s)  Cu(aq) + 2e - E° = -0.34 V The more positive potential is the reduction potential. The more positive potential is the reduction potential.

36. Maximum Standard Cell Voltage (E° cell ) E° cell = E° reduced + E° oxidized Example: (cathode) (anode) Cu +2 (aq) + 2e -  Cu(s)+0.34 V Zn +2 (aq) + 2e -  Zn(s) -0.76 V * Cu has a higher voltage and is therefore the reduction reaction so we need to change the sign of the zinc reduction to make it oxidation. E° = E° Cu +2 /Cu + E° Zn/Zn +2 = (+0.34 V )+ (+0.76 V) = +1.10 V

37. Electrons will spontaneously flow in the direction that gives a positive cell potential. Electrons will spontaneously flow in the direction that gives a positive cell potential. Example 1: Sn +2 (aq) + 2e -  Sn(s)-0.14 V Ag + (aq) + e -  Ag(s)+0.80 V Example 2: Fe(s) + Pb +2 (aq)  Fe +2 (aq) + Pb(s)

38. Electolytic Cell The REDOX reaction in an electrolytic cell is non-spontaneous, and therefore electrical energy (battery) is required to induce a chemical reaction The REDOX reaction in an electrolytic cell is non-spontaneous, and therefore electrical energy (battery) is required to induce a chemical reaction (electrical energy  chemical energy)

39. Battery Anode (+) Cathode (-) NaCl  Na + + Cl -

40. Electrolytic Cell Electrolysis: the process in which electricity breaks down a compound. Example: NaCl Electrolysis: the process in which electricity breaks down a compound. Example: NaCl The negative end of the battery is attached to the cathode (reduction) which makes it negatively charged; the positive end of the battery is attached to the anode (oxidation) which makes it positively charged. The negative end of the battery is attached to the cathode (reduction) which makes it negatively charged; the positive end of the battery is attached to the anode (oxidation) which makes it positively charged. Positive ions (Na + ) in solution are attracted to the cathode which is negatively charged. Positive ions (Na + ) in solution are attracted to the cathode which is negatively charged. Negative ions (Cl - ) in solution are attracted to the anode which is positively charged. Negative ions (Cl - ) in solution are attracted to the anode which is positively charged.

41. CATHODE (reduction): Na + + e -  Na 0 We have created neutral sodium which is never found in nature (so reactive). We have created neutral sodium which is never found in nature (so reactive). ANODE (oxidation) Cl -  e - + Cl 0 We have created neutral chlorine which is never found in nature (so reactive) We have created neutral chlorine which is never found in nature (so reactive)

42. Voltaic Electrolytic Voltaic Electrolytic Spontaneous Spontaneous Chemical  Electrical Energy Chemical  Electrical Energy Anode – oxidation (-) Anode – oxidation (-) Cathode – reduction (+) Cathode – reduction (+) Example: BATTERY Example: BATTERY Electrons travel from anode to cathode (wire) Electrons travel from anode to cathode (wire) More reactive metal ALWAYS the site of oxidation More reactive metal ALWAYS the site of oxidation Salt bridge is for the flow of ions from one half-cell to another Salt bridge is for the flow of ions from one half-cell to another Non-spontaneous Electrical  Chemical Requires a battery as a source of energy Anode – oxidation (+) Cathode – reduction (-) (+) charged ion moves toward the cathode (-) charged ion moves toward the anode

43. Electroplating Another example of an electrolytic cell. Another example of an electrolytic cell. The process of electroplating requires a layer of metal such as silver or copper, coating or covering any object to be plated (spoon or fork) The process of electroplating requires a layer of metal such as silver or copper, coating or covering any object to be plated (spoon or fork) The item being plated is the cathode - reduction/(-) The item being plated is the cathode - reduction/(-) The electrode must be the same metal that you are plating the object in. The electrode must be the same metal that you are plating the object in. Cathode (-)

44. Review Questions 1) In which substance does chlorine have an oxidation number of +1? A) HClO 2 B) HClO C) Cl 2 D) HCl 2) What is the oxidation state of nitrogen in NaNO 2 ? A) +1 B) +2 C) +3 D) +4

45. 3) Which balanced equation represents a redox reaction? A) AgNO 3 + NaCl  AgCl + NaNO 3 B) BaCl 2 + K 2 CO 3  BaCO 3 + 2KCl C) CuO + CO  Cu + CO 2 D) HCl + KOH  KCl + H 2 O 4) Given the reaction: Mg(s) + 2H+(aq) + 2Cl - (aq)  Mg +2 (aq) + 2Cl - (aq) + H 2 (g) Which species undergoes oxidation? A) Cl - (aq) B) Mg(s) C) H + (aq) D) H 2 (g)

46. 5) Given the reaction: Zn(s) + 2HCl(aq)  ZnCl2(aq) + H2(g) Which statement correctly describes what occurs when this reaction takes place in a closed system? A) There is a net gain of mass. B) Atoms of Zn(s) lose electrons and are oxidized. C) Atoms of Zn(s) gain electrons and are reduced. D) There is a net loss of mass. 6) Given the reaction that occurs in an electrochemical cell: Zn(s) + CuSO 4 (aq)  ZnSO 4 (aq) + Cu(s) During this reaction, the oxidation number of Zn changes from A) 0 to +2 B) 0 to -2 C) +2 to 0 D) -2 to 0

47. 7) Given the unbalanced equation: __Br 2 + __Sn  __Br - + __Sn +2 When the equation is correctly balanced using the smallest whole-number coefficients, the coefficient of Br - is A)1 B) 2 C)3 D) 4 8) What is the reducing agent in the reaction: Pb + 2AgNO 3  Pb(NO 3 ) 2 + 2Ag? A) Pb B) NO -3 C) Ag + D) Ag 9) According to the Activity Series chemistry reference table, which molecule is most easily reduced? A) I 2 B) Br 2 C) Cl 2 D) F 2

48. 10) Given the reaction: 4Al(s) + 3O 2 (g)  2Al 2 O 3 (s) (a) Write the balanced oxidation half- reaction for this oxidation-reduction reaction. (b) What is the oxidation number of oxygen in Al 2 O 3 ?

49. 11) The diagram below represents a voltaic cell at 298 K and 1 atmosphere. When the switch is closed, electrons flow from A) Mg(s) to Ag(s) B) Ag(s) to Mg(s) C) Ag + (aq) to Mg +2 (aq) D) Mg +2 (aq) to Ag + (aq)

50. Questions 12 and 13 refer to the following: The diagram below shows a spoon that will be electroplated with nickel metal. 12) Does the chemical cell diagram represent a voltaic or an electrolytic cell? [Give one reason to support your answer.] 13) Does the spoon represent the anode or the cathode in this electrochemical cell? [Give one reason to support your answer.]