Practice Packet Unit 13: Electrochemistry (RedOx)

Regents Chemistry: Mr. Palermo Practice Packet Unit 13: Electrochemistry (RedOx) Redox and Batteries? Ain t nobody got time for that!!! 1

Lesson 1: Oxidation States Oxidation numbers are very important in this chapter Redox Reactions. Without the complete understanding of how to assign these numbers, we cannot move ahead with this chapter. They are much like ionic charges, except the every element will be assigned a number. The most important rules that cannot be broken are: Free elements are zero Group 1 is +1 Group 2 is +2 Fluorine is -1 The sum of oxidation states in a compound is 0. The sum of oxidation states in a polyatomic ion is equal to the charge of that ion. Assign oxidation numbers to each element in the following (use the Periodic Table to help you) (a) NaCl Na Cl (b) H 2 S H S (c) H 2 O H O (d) CO 2 C O (e) H 2 SO 4 H S O (f) FeCO 3 Fe C O (g) AgI Ag I (h) H 2 H (i) PbCl 2 Pb Cl (j) BaCO 3 Ba C O (k) Fe 2 O 3 Fe O (l) I 2 I (m) BeO Be O (n) CaF 2 Ca F (o) FeCl 3 Fe Cl (p) PF 5 P F (q) H 3 PO 4 H P O (r) KCl K Cl (s) K 2 O K O (t) O 3 O (u) LiH Li H (v) HBr H Br (w) Li + Li 3- (x) PO 4 P O What is the oxidation number of each element in the following reactions: 1. AgNO 3 + FeCl 3 AgCl + Fe(NO 3 ) 3 2. MgSO 4 + Ca(OH) 2 Mg(OH) 2 + CaSO 4 3. FeBr 2 + Br 2 FeBr 3 2

Lesson 2: Identifying a Redox Reaction A redox reaction is a reaction in which electrons are transferred from one element to another. The reaction involves at least two elements, one that will give up an electron, and one that will receive that electron. The term redox comes from two words, oxidation and reduction. If something is oxidized, it burns in oxygen, as shown below: Mg + O Mg +2 + O -2 Write out the Lewis dot diagrams for each species above. As you can see, Mg is being oxidized as it loses its two valence electrons, while oxygen gains two. LEO = GER = Indicate which element is being oxidized, and which is being reduced. 1. Cr 3+ + Fe 2+ Cr 2+ + Fe 3+ 2. F 2 + O 2- F 1- + O 2 3. Sn + N 5+ Sn 4+ + N 4+ 4. NaCl Na + + Cl - 5. Cu 2 O Cu + O 2 6. Cl 2 + KBr KCl + Br 2 7. CH 4 + O 2 CO 2 + H 2 O 8. H 3 PO 4 + Ca(OH) 2 Ca 3 (PO 4 ) 2 + H 2 O 1. 2. 3. 4. 5. 6. 7. 8. Oxidized Reduced 3

Redox reactions are usually synthesis reactions, decomposition, combustion or single replacement reactions. Double replacement and neutralization reactions are NOT redox reactions. Usually they are easy to spot because if an element goes from being free (with an oxidation number of 0) to being in a compound (with a new oxidation number) it shows there was an exchange of electrons. In the following examples, identify what type of reaction they are and then state if they are redox reactions. 1. N 2 + O 2 2NO 2. Cl 2 + 2NaBr NaCl + Br 2 3. 2NaOH + HCl H 2 O + NaCl Determine if the following reactions are REDOX reactions. Then determine which species is oxidized and which is reduced. 1) NaOH + HCl NaCl + H 2 O 2) N 2 + 3H 2 2NH 3 3) NaCl + AgNO 3 NaNO 3 + AgCl 4) Zn + CuSO 4 ZnSO 4 + Cu 5) 2KCl + Br 2 2KBr + Cl 2 6) 2H 2 + O 2 2H 2 O 7) Ba(OH) 2 + H 2 SO 4 BaSO 4 + 2H 2 O 8) 2Mg + O 2 2MgO 9) 2AgNO 3 + CuSO 4 Ag 2 SO 4 + Cu(NO 3 ) 2 What type of reaction (S, D, C, SR, or DR) is NEVER redox? 4

Spontaneous Reactions and Activity Series Use Table J to help you with the following questions. 1. Write the oxidation and reduction half reaction for: Ca + Cu 2+ Ca 2+ + Cu 2. According to Table J, the element higher on the list will oxidize. Which element is oxidizing here and does that mean that this reaction is spontaneous? 3. Write the oxidation and reduction half reaction for Mg + Ca 2+ Mg 2+ + Ca 4. Which element is oxidizing in number 3 and is the reaction spontaneous? 5. Which of the following ions is most easily oxidized? a. F- b. Cl- c. Br- d. I- 6. Which element is more easily reduced? a. Cu b. Mg c. Al d. Zn 7. Which element will reduce Mg 2- to Mg? a. Fe b. Ba c. Pb d. Ag 8. Which ion will oxidize Fe? a. Zn 2+ b. Ca 2+ c. Mg 2+ d. Cu 2+ 9. Which metal will react spontaneously with Ag + but not Zn 2+? a. Cu b. Au c. Al d. Mg 10. Which reaction will take place spontaneously? a. Mg + Ca 2+ Mg 2+ + Ca b. Ba + 2Na + Ba 2+ + 2Na c. Cl 2 + 2F - 2Cl - +F 2 d. I 2 + 2Br - 2I - + Br 2 5

Lesson 3: Half Reactions A half reaction shows either the oxidation or reduction portion of a redox equation including if the electrons are gained or lost. A reduction half reaction shows an atom or ion gaining one of more e - : Fe 3+ + 3e - Fe Notice that the e - is one the left. An oxidation half reaction shows an atom or ion losing one or more e-: Mg Mg 2+ + 2 e - Notice that the e - is one the right. In a half reaction, only on element is shown and the charges must be conserved. To write a half reaction, first assign all the oxidation numbers to all the elements. Second, cross out any elements that are spectators (they do not change oxidation number). Then, write a half reaction showing a change in oxidation state and label which element is being oxidized and which is being reduced. Last, add in the number of electrons needed to conserve the charge. 1. Cr 3+ + Fe 2+ Cr 2+ + Fe 3+ 2. F 2 + O 2- F 1- + O 2 3. Sn + N 5+ Sn 4+ + N 4+ 4. NaCl Na + + Cl - For each reaction, decide if it is redox. If it is redox, write the half reactions below. If they are not redox, leave the answers blank. 5. Cu 2 O Cu + O 2 6. Cl 2 + KBr KCl + Br 2 7. CH 4 + O 2 CO 2 + H 2 O 8. H 3 PO 4 + Ca(OH) 2 Ca 3 (PO 4 ) 2 + H 2 O Oxidation Reduction 1. 2. 3. 4. 5. 6. 7. 8. 6

Fill in the table below. 1 H 2 + O 2 H 2 O Which element is reducing? Which element is oxidizing? Reducing half Reaction Oxidizing half reaction How many e- are transferred? 2 Fe + Zn 2+ Fe 2+ + Zn 3 2Al + 3Fe 2+ 2Al +3 + 3Fe 4 Cu + 2AgNO 3 Cu(NO 3 ) 2 + 2Ag Balancing Redox Equations Practice: Balance the following redox reactions 1. CsF + Na NaF + Cs 2. NaCl + Br 2 NaBr + Cl 2 7

3. MgCl 2 + Cr Mg + CrCl 3 4. Pb + AgNO 2 Pb(NO 2 ) 2 + Ag 5. Fe +2 + Cu + Fe +3 + Cu 6. Cr + CuBr 2 CrBr 3 + Cu 7. Zn + CuO ZnO + Cu 8

RedOx Regents Review 1. Which reaction is an example of an oxidationreduction reaction? (1) AgNO 3 + KI AgI + KNO 3 (2) Cu + 2AgNO 3 Cu(NO 3 ) 2 + 2Ag (3) 2KOH + H 2 SO 4 K 2 SO 4 + 2H 2 O (4) Ba(OH) 2 + 2HCl BaCl 2 + 2H 2 O 2. In an oxidation-reduction reaction, reduction is defined as the (1) loss of protons (3) loss of electrons (2) gain of protons (4) gain of electrons 3. When a lithium atom forms a Li + ion, the lithium atom (1) gains a proton (3) loses a proton (2) gains an electron (4) loses an electron 4. Which type of reaction occurs when nonmetal atoms become negative nonmetal ions? (1) oxidation (3) substitution (2) reduction (4) condensation 5. When a neutral atom undergoes oxidation, the atom s oxidation state (1) decreases as it gains electrons (2) decreases as it loses electrons (3) increases as it gains electrons (4) increases as it loses electrons 6. In a redox reaction, there is a conservation of (1) mass, only (2) both mass and charge (3) neither mass nor charge 7. In any redox reaction, the substance that undergoes reduction will (1) lose e- & have a decrease in oxidation number (2) lose e- & have an increase in oxidation number (3) gain e- & have a decrease in oxidation number (4) gain e- & have an increase in oxidation number 8. What occurs during the reaction below? (1) The manganese is reduced and its oxidation number changes from +4 to +2. (2) The manganese is oxidized and its oxidation number changes from +4 to +2. (3) The manganese is reduced and its oxidation number changes from +2 to +4. (4) The manganese is oxidized and its oxidation number changes from +2 to +4. 9. Given the balanced equation: What is the total number of moles of electrons lost by 2 moles of Al(s)? (1) 1 mole (3) 3 moles (2) 6 moles (4) 9 moles 10.Given the balanced equation: Mg (s) + Ni 2+ (aq) Mg 2+ (aq) + Ni (s) What is the total number of moles of electrons lost by 2 moles of Mg(s)? (1) 1.0 mol (3) 3.0 mol (2) 2.0 mol (4) 4.0 mol 11. Given the equation representing a reaction: Mg (s) + Ni 2+ (aq) Mg 2+ (aq) + Ni (s) What is the total number of moles of e- lost by Mg when 2.0 moles of e- are gained by Ni 2+ (aq)? (1) 1.0 mol (3) 3.0 mol (2) 2.0 mol (4) 4.0 mol 12. Given the reaction: Which species undergoes oxidation? (1) Mg(s) (3) Cl (aq) (2) H + (aq) (4) H 2 (g) 9

13. Given the redox reaction: As the reaction takes place, there is a transfer of (1) electrons from Al to Cr 3+ (2) electrons from Cr 3+ to Al (3) protons from Al to Cr 3+ (4) protons from Cr 3+ to Al 14. Given the redox reaction: 2 Fe 3+ + 3 Zn 2 Fe + 3 Zn 2+ As the reaction takes place, there is a transfer of electrons (1) from Fe 3+ to Zn (2) from Zn to Fe 3+ (3) from Zn 2+ to Fe (4) from Fe to Zn 2+ 14. Circle the electrons in the half-reactions below and identify as oxidation or reduction. (a) Br2 + 2 e 2 Br (d) Cl2 + 2 e 2 Cl (g) Cu 2+ + 2 e Cu (b) Na Na + + e (e) Na + + e Na (h) Fe Fe 2+ + 2 e (c) Ca 2+ + 2e Ca (f) S 2 S + 2e (i) Mn 7+ + 3 e Mn 4+ 15. Complete the half-reactions below by ADDING in electrons to the correct side in order to equalize charge (show conservation of charge). (a) Fe 2+ Fe 3+ (b) K K + (c) Sn 4+ Sn 2+ (d) Cr 6+ Cr 3+ (e) O 2 2O (f) Mn 3+ Mn 4+ (g) Cr 2+ Cr 3+ (h) Cl 7+ Cl 1+ (i) 3Cl 2 6Cl (j) 4H + 2H 2 16. For each of the following equations, write the reduction half-reaction and the oxidation half-reaction. Reduction: (a) Cl 2 + 2 KBr 2 KCl + Br 2 Oxidation: (b) Cu + 2 Ag + 2 Ag + Cu 2+ Reduction: Oxidation: Reduction: (c) 2 Mg + O 2 2 MgO (d) 2 F 2 + 2 H 2 O 4 HF + O 2 17. Which half-reaction correctly represents oxidation? (1) Fe(s) Fe 2+ (aq) + 2e Oxidation: Reduction: Oxidation: (2) Fe 2+ (aq) Fe(s) + 2e (3) Fe(s) + 2e Fe 2+ (aq) (4) Fe 2+ (aq) + 2e Fe(s) 10

18. Which equation shows a conservation of both mass and charge? (1) Cl 2 + Br Cl + Br 2 (2) Cu + 2Ag + Cu 2+ + Ag + (3) Zn + Cr 3+ Zn 2+ + Cr (4) Ni + Pb 2+ Ni 2+ + Pb 19. Given the balanced ionic equation: 20. Which half-reaction equation represents the reduction of a potassium ion? (1) K + + e K (3) K + K + e (2) K + e K + (4) K K + + e 21. Given the equation: Which equation represents the oxidation half-reaction? (1) Zn(s) + 2e Zn 2+ (aq) (2) Zn(s) Zn 2+ (aq) + 2e (3) Cu 2+ (aq) Cu(s) + 2e (4) Cu 2+ (aq) + 2e Cu(s) The reduction half-reaction is (1) Al Al 3+ + 3e (2) Cu 2+ + 2e Cu (3) Al + 3e Al 3+ (4) Cu 2+ Cu + 2e 22. Base your answers to the questions below on the following redox reaction, which occurs in a battery. Zn + Cr 3+ Zn 2+ + Cr (a) Write the half-reaction for the reduction that occurs. (b) Write the half-reaction for the oxidation that occurs. (c) Balance the equation below using the smallest whole-number coefficients. Zn + Cr 3+ Zn 2+ + Cr (d) Which species loses electrons and which species gains electrons? (e) State what happens to the number of protons in a Zn atom when it changes to Zn 2+ as the redox reaction occurs. 23. The outer structure of the Statue of Liberty is made of copper metal. The framework is made of iron. Over time, a thin green layer (patina) forms on the copper surface. (a) When copper oxidized to form this patina layer, the copper atoms became copper(ii) ions (Cu 2+ ). Write a balanced half-reaction for this oxidation of copper. Cu Cu 2+ + 2e (b) Where the iron framework came in contact with the copper surface, a reaction occurred in which iron was oxidized. Using information from Reference Table J, explain why the iron was oxidized. Iron is a more active metal. 11

24. Litharge, PbO, is an ore that can be roasted (heated) in the presence of carbon monoxide, CO, to produce elemental lead. The reaction that takes place during this roasting process is represented by the balanced equation below. PbO(s) + CO(g) Pb(l) + CO 2 (g) (a) Write the balanced equation for the reduction half-reaction that occurs during this roasting process. (b) Determine the oxidation number of carbon in carbon monoxide. 25. The catalytic converter in an automobile changes harmful gases produced during fuel combustion to less harmful exhaust gases. In the catalytic converter, nitrogen dioxide reacts with carbon monoxide to produce nitrogen and carbon dioxide. In addition, some carbon monoxide reacts with oxygen, producing carbon dioxide in the converter. These reactions are represented by the balanced equations below. Determine the oxidation number of carbon in each carbon compound in reaction 2. Your response must include both the sign and value of each oxidation number. +2 for carbon in CO and +4 for carbon in CO 2. 26. A flashlight can be powered by a rechargeable nickel-cadmium battery. The unbalanced equation below represents the reaction that occurs as the battery produces electricity. When a nickel-cadmium battery is recharged, the reverse reaction occurs. (a) Balance the equation above using the smallest whole-number coefficients. 1,1,2,1,1 (b) Determine the change in oxidation number for Cd. from 0 to +2 (c) Explain why Cd would be above Ni if placed on Table J. Cd is more reactive than Ni. 12

Lesson 4: Electrochemical Cells Voltaic Cells Directions: In each of the following, determine which element oxidized easier on table J. Then label the anode, cathode, direction of e- flow, and the half reactions. 1. 2. 3. 4. Additional Questions: 4. On diagram 1, which way will anions travel through the salt bridge? 5. On diagram 2, towards which electrode will cations travel through the salt bridge? 6. On diagram 3, how many e- are exchanged per mole of Mg? 7. On diagram 4, how many e- are transferred between Ag and Ni? 8. On all diagrams, at which electrode does oxidation occur? 9. On all diagrams, at which electrode does reduction occur? 10. On all diagrams, from which electrode will electrons travel? 11. What is the purpose of the salt bridge? 13

Lesson 5: Electrolytic Cells 1. Which element is oxidizing in the diagram to the left? 2. According to Table J, should that element oxidize? 3. Is this reaction spontaneous? 4. What type of cell is it? 5. Label the anode and the cathode with charges on the diagram to the left. 6. Write the overall reaction for this cell. 7. Water is being decomposed using a battery in the diagram to the right. Write the equation for the decomposition of water. 8. Which element is being oxidized? How many e- are lost? 9. According to Table J, should that element oxidize? 10. Is this reaction spontaneous? 11. What type of cell is it? 12. Label where oxidation and reduction on the diagram. 13. Why is more H 2 gas being formed that O 2 gas? 14. Only one element is being used in the diagram to the left. The silver plate ionizes and the ions attach to the spoon. Show a half reaction for silver oxidizing. 15. Label the anode and cathode with charges on the diagram. 16. Show the direction of e- flow through the wire. 17. Is this reaction spontaneous? How can you tell? 14

18. Write the half reaction for the anode in the diagram to the left: 19. Write the half reaction for the cathode (use Fe +2 ): 20. How many e- are transferred per mole? 21. Show the direction of e- flow through the power supply. 22. Why is a power source needed in this cell? 23. Show a half reaction for silver reducing. 24. Label the cathode on the diagram to the right. 25. What will happen to the mass of the key? 26. What will happen to the mass of the silver metal? 27. Show the direction of e- flow through the wire on the diagram to the right. 29. Why will the mass of the key increase? 30. Label the anode, cathode, and the direction of e-flow through the wire. 31. State the difference between voltaic and electrolytic cells in terms of spontaneity. 32. State the difference between voltaic and electrolytic cells in terms of energy being released or absorbed. 15

Electrochemical Cells Regents Questions 1. In a voltaic cell, chemical energy is converted to (1) electrical energy, spontaneously (2) electrical energy, nonspontaneously (3) nuclear energy, spontaneously (4) nuclear energy, nonspontaneously 2. A voltaic cell spontaneously converts (1) electrical energy to chemical energy (2) chemical energy to electrical energy (3) electrical energy to nuclear energy (4) nuclear energy to electrical energy 3. A voltaic cell differs from an electrolytic cell in that in a voltaic cell (1) energy is produced when the reaction occurs (2) energy is required for the reaction to occur (3) both oxidation and reduction occur (4) neither oxidation nor reduction occurs 4. Which half-reaction can occur at the anode in a voltaic cell? (1) Ni 2+ + 2e- Ni (2) Sn + 2e- Sn 2+ (3) Zn Zn 2+ + 2e- (4) Fe 3+ Fe 2+ + e- 5. Which process requires an external power source? (1) neutralization (3) fermentation (2) synthesis (4) electrolysis 6. Which energy transformation occurs when an electrolytic cell is in operation? (1) chemical energy electrical energy (2) electrical energy chemical energy (3) light energy heat energy (4) light energy chemical energy 7. What is the purpose of the salt bridge in a voltaic cell? (1) It blocks the flow of electrons. (2) It blocks the flow of positive and negative ions. (3) It is a path for the flow of electrons. (4) It is a path for the flow of positive and negative ions. 8. Which statement is true for any electrochemical cell? (1) Oxidation occurs at the anode, only. (2) Reduction occurs at the anode, only. (3) Oxidation occurs at both the anode and the cathode. (4) Reduction occurs at both the anode and the cathode. 9. Given the balanced equation representing a reaction occurring in an electrolytic cell: 2NaCl (l) 2Na (l) + Cl 2(g) Where is Na(l) produced in the cell? (1) at the anode, where oxidation occurs (2) at the anode, where reduction occurs (3) at the cathode, where oxidation occurs (4) at the cathode, where reduction occurs Answer questions 10 and 11 using the diagram below, which represents an electrochemical cell. 10. When the switch is closed, in which half-cell does oxidation occur? 11. What occurs when the switch is closed? (1) Zn is reduced. (2) Cu is oxidized. (3) Electrons flow from Cu to Zn. (4) Electrons flow from Zn to Cu. 16

Use the diagram of a key being plated with copper to answer questions 12 through 15. 12. What is the name of the process shown in the diagram? 13. What is the purpose of the battery in this electrolytic cell? 14. Which electrode, A or B, attracts positive copper ions? 15. Given the reduction reaction for this cell: Cu 2+ (aq) + 2e Cu(s) This reduction occurs at (1) A, which is the anode (3) B, which is the anode (2) A, which is the cathode (4) B, which is the cathode 16. Aluminum is one of the most abundant metals in Earth s crust. The aluminum compound found in bauxite ore is Al 2 O 3. Over one hundred years ago, it was difficult and expensive to isolate aluminum from bauxite ore. In 1886, a brother and sister team, Charles and Julia Hall, found that molten (melted) cryolite, Na 3 AlF 6, would dissolve bauxite ore. Electrolysis of the resulting mixture caused the aluminum ions in the Al 2 O 3 to be reduced to molten aluminum metal. This less expensive process is known as the Hall process. (a) Write the oxidation state for each of the elements in cryolite. Na: +1 Al: +3 F: 1 (b) Write the balanced half-reaction equation for the reduction of Al 3+ to Al. Al 3+ + 3e Al (c) Explain, in terms of ions, why molten cryolite conducts electricity. There are freely moving ions in the molten c (d) Explain, in terms of electrical energy, how the operation of a voltaic cell differs from the operation of an electrolytic cell used in the Hall process. Include both the voltaic cell and the electrolytic cell in your answer. Electrolysis uses electrical energy. Voltaic cells produce electrical energy. 17. Base your answers to the following questions on the diagram of the voltaic cell below. (a) Identify the anode and the cathode. Anode = Pb Cathode = Ag (b) Write the oxidation and reduction half-reactions for this voltaic cell. Red: 2Ag + + 2 e- 2Ag 2 e- (c) What is the total number of moles of electrons needed to completely reduce 6 moles of Ag + (aq) ions? (d) Describe the direction of electron flow between the electrodes. Electrons flow from the Pb electrode to the Ag electrode 17

(e) State the purpose of the salt bridge in this cell. Maintains a balance of charge; allows ions to migrate (f) State the electrode to which positive ions migrate when the switch is closed. Ag electrode (the cathode) to balanced the electrons that are arriving (g) As this voltaic cell operates, the mass of the Ag(s) electrode increases. Explain, in terms of silver ions and silver atoms, why this increase in mass occurs. At the Ag electrode, silver ions in the water are being reduced to solid silver atoms. The solid silver deposits on the Ag electrode, which increases its mass. 18. Base your answers to the following questions on the diagram below, which represents a voltaic cell at 298 K and 1 atm. (a) In which half-cell will oxidation occur when switch S is closed? Zn, anode, right (b) Write the balanced half-reaction equation that will occur in half-cell 1 when switch S is closed. Pb 2+ + 2e (c) Describe the direction of electron flow between the electrodes when switch S is closed. anode to cathode 19. Base your answers to the following questions on the information below. Underground iron pipes in contact with moist soil are likely to corrode. This corrosion can be prevented by applying the principles of electrochemistry. Connecting an iron pipe to a magnesium block with a wire creates an electrochemical cell. The magnesium block acts as the anode and the iron pipe acts as the cathode. A diagram of this system is shown below. (a) State the direction of the flow of electrons between the electrodes in this cell. From Mg (anode) to Fe (cathode) (b) Explain, in terms of reactivity, why magnesium is preferred over zinc to protect underground iron pipes. Your response must include both magnesium and zinc. ore readily than Zn. The salt bridge allows ions to flow between the half-cells. Maintains a balance of charge 18

20. Base your answers to the following questions on the diagram and balanced equation below, which represent the electrolysis of molten NaCl. (a) When the switch is closed, which electrode will attract the sodium ions? cathode, one on the right (b) What is the purpose of the battery in this electrolytic cell? energy source (c) Write the balanced half-reaction for the reduction that occurs in this electrolytic cell. Na + + e Na OR 2 21. State two similarities and two differences between voltaic and electrolytic cells. Electrons flow anode to cathode Diff: voltaic: produces energy, spongy Electrolytic: requires energy, nonspontaneous, 1 cell, electrical energy chemical energy 22. The apparatus shown in the diagram consists of two inert platinum electrodes immersed in water. A small amount of an electrolyte, H 2 SO 4, must be added to the water for the reaction to take place. The electrodes are connected to a source that supplies electricity. (a) What type of electrochemical cell is shown? electrolytic or electrolysis (b) What particles are provided by the electrolyte that allow an electric current to flow? Ions, charged particles, H 3 O +, or SO 4 2 23. The diagram below shows a system in which water is being decomposed into oxygen gas and hydrogen gas. Litmus is used as an indicator in the water. The litmus turns red in test tube 1 and blue in test tube 2. The oxidation and reduction occurring in the test tubes are represented by the balanced equations below. (a) Identify the information in the diagram that indicates this system is an electrolytic cell. A battery is part of the cell and is providing energy that causes the reaction. Electricity is used to operate the cell. 19

(b) Determine the change in oxidation number of oxygen during the reaction in test tube 1. -2 to 0 (c) Explain, in terms of the products formed in test tube 2, why litmus turns blue in test tube 2. Litmus turns blue when hydroxide ions are produce 25. Which reaction occurs spontaneously? (1) Cl 2(g) + 2NaBr (aq) Br 2(l) + 2NaCl (aq) (2) Cl 2(g) + 2NaF (aq) F 2(g) + 2NaCl (aq) (3) I 2(s) + 2NaBr (aq) Br 2(l) + 2NaI (aq) (4) I 2(s) + 2NaF (aq) F 2(g) + 2NaI (aq) 26. Which metal reacts spontaneously with a solution containing zinc ions? (1) magnesium (3) copper (2) nickel (4) silver 27. Which metal with react with Zn 2+ spontaneously, but will not react with Mg 2+? (1) Al (3) Ni (2) Cu (4) Ba 28. Which of the following metals has the least tendency to undergo oxidation? (1) Ag (3) Zn (2) Pb (4) Li 29. Because tap water is slightly acidic, water pipes made of iron corrode over time, as shown by the balanced ionic equation: 2Fe(s) + 6H + (aq) 2Fe 3+ (aq) + 3H 2 (g) Explain, in terms of chemical reactivity, why copper pipes are less likely to corrode than iron pipes. 30. Identify one metal that does not react spontaneously with HCl(aq). 20