Complete throughout unit. Due on test day!

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1 Name Unit 8: REDOX and Electrochemistry Skills: 1. Assigning Oxidation Numbers 2. Identifying Oxidation and Reduction 3. Writing Half Reactions 4. Balance Redox Reaction (Flipped) Unit 8: Vocabulary: Word Reduction Period 5. Table J and Spontaneous Reactions 6. Distinguish Between electrochemical batteries: Voltaic and Electrolytic 7. Solve Voltaic Cell Problem Complete throughout unit. Due on test day! Definition Oxidation Half Reaction Anode Cathode Battery Power Source Voltage Voltaic Cell Electrolytic Cell Electrolysis Electroplating Salt Bridge Spontaneous

2) Make observations! 3) Filter you product in the class filter at the front! 4) What does the Flame Test tell us about our product? Complete and Balance!

2 What is REDOX?... REDuction-OXidation Reactions! Reactions that involve the TRANSFER OF ELECTRONS; both reduction and oxidation must happen! Reduction = by an atom or ion; goes / Oxidation = LOSS OF ELECTRONS by an atom or ion; OXIDATION NUMBER goes /. Green Chunks Experiment! 1) Put your Aluminum Foil into your Copper Chloride solution! 2) Make observations! 3) Filter you product in the class filter at the front! 4) What does the Flame Test tell us about our product? Complete and Balance! Al + CuCl 2 à + Aluminum is above Cu on Table J so it will replace it! Notice how Al is all by itself (on left of arrow) with a zero charge and then bonded (on right of arrow) where it takes on a charge! Skill 1: Assigning Oxidation Numbers Oxidation Numbers = Charge numbers o UNCOMBINED ELEMENTS (elements not bonded to another element) have an oxidation number of. Ex: Mg: O2: Cu: Fe: o In, the sum of the CHARGES for all elements must ADD UP TO ZERO. Ex: NaCl H2SO4 First element should always be positive (+) and second element should always be negative (-).

1. Assign oxidation numbers to each element in the following (use the PTable to help you) (a) NaCl Na Cl (f) FeCO3 Fe C O (b) H2S H S (g) AgI Ag I (c) H2O H O (h) H2 H (d) CO2 C O (i) PbCl2 Pb Cl (e)

3 1. Assign oxidation numbers to each element in the following (use the PTable to help you) (a) NaCl Na Cl (f) FeCO3 Fe C O (b) H2S H S (g) AgI Ag I (c) H2O H O (h) H2 H (d) CO2 C O (i) PbCl2 Pb Cl (e) H2SO4 H S O (j) BaCO3 Ba C O Assign oxidation numbers to each element in the polyatomic ions: k) (SO4) - 2 S O l) (OH) - O H SKILL #2: Identifying Oxidation & Reduction - o A reaction is REDOX if...oxidation NUMBERS. Trick 1: SINGLE REPLACEMENT REACTIONS are always REDOX! Example: Zn + HCl à + * / are bonded on one side and alone on the other Trick 2: DOUBLE REPLACEMENT REACTIONS are NOT REDOX! Example: Na(OH) + HCl à + *Charges stay the same for all elements in the rxn

4 FOR 2 ELEMENTS WITHIN A REACTION: o Reduction (GER) = OF ELECTRONS by an atom or ion; oxidation number goes down/reduces o Oxidation (LEO)= OF ELECTRONS by a atom or ion; oxidation number goes up/oxidizes C + H2O à CO + H2 Charge: Increase/ Decrease Electrons: Lost/ Gained Oxidized or Reduced? C 0 H +1 MnO2 + 4HCl à MnCl2 + 2H2O Charge: Increase/ Decrease Electrons: Lost/ Gained Oxidized or Reduced? Cl -1 Mn +4 FeBr2 + Br2 à FeBr3 Charge: Increase/ Decrease Electrons: Lost/ Gained Oxidized or Reduced? Cl -1 Mn +4 Answer: 1. In any redox reaction, the substance that undergoes reduction will (1) lose electrons and have a decrease in oxidation number (2) lose electrons and have an increase in oxidation number (3) gain electrons and have a decrease in oxidation number (4) gain electrons and have an increase in oxidation number 2. 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 electron

5 Name Period For each of the following: a) Assign oxidation numbers for each element/species; b) Indicate the element/species, if any, that are oxidized and reduced. 1) Zn + Pb(NO3)2à Pb + Zn(NO3)2 oxidized substance reduced substance 2) FeCO3 + 3COà Fe + 3CO2 oxidized substance reduced substance 3) 3H2 + N2à NH3 oxidized substance reduced substance Complete the following Question: 1. Which reaction is an example of an oxidation-reduction reaction? (1) AgNO3 + KI à AgI + KNO3 (2) Cu + 2AgNO3 à Cu(NO3)2 + 2Ag (3) 2KOH + H2SO4 à K2SO4 + 2H2O (4) Ba(OH)2 + 2HCl à BaCl2 + 2H2O 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 2. In an oxidation-reduction reaction, reduction is defined as the (1) loss of protons (2) loss of electrons (3) gain of protons (4) gain of electron 4. Which type of reaction occurs when nonmetal atoms become negative nonmetal ions? (1) oxidation (3) substitution (2) reduction (4) condensation

6 SKILL #3: Writing Half-Reactions o Half reactions allow us to show the of in a redox rxn. For each redox reaction, we can illustrate two REACTIONS. One half-reaction shows and other shows REDUCTION. Example of a Reduction Half Reaction: Fe e- àfe Electrons on left hand side, GAINED in the rxn (GER). Notice also how the charge for Fe goes down from left to right, REDUCTION (GER). Charge goes down because Fe GAINED e -. Example of an Oxidation Half Reaction: Fe àfe e- Electrons on left hand side, LOST in the rxn (LEO). Notice also how the charge for Fe goes up from left to right, OXIDATION (LEO). Charge goes up because Fe LOST e -. o Always ADD ELECTRONS to the side of rxn that has a HIGHER TOTAL CHARGE (Remember: electrons are ) FOLLOWING THE LAW OF CONSERVATION: o Half reactions follow the LAW OF of. This means that there must be the SAME NUMBER OF ATOMS on both sides of the reaction arrow. o There must also be a CONSERVATION OF. In half reactions, the NET CHARGE MUST BE THE ON BOTH SIDES of the equation, although it doesn t necessarily need to equal zero. Example 1: Mg + ZnCl 2 à MgCl 2 + Zn o OXIDATION Half Reaction: o REDUCTION Half Reaction: Example 2: (balance masses): Hg + I 2 à HgI o OXIDATION Half Reaction: o REDUCTION Half Reaction:

7 Circle the electrons in the half-reactions below and identify as oxidation or reduction (a) Br2 + 2 e à2 Br (d) Cl2 + 2 e à 2 Cl (g) Cu e à Cu (b) Na à Na + + e (e) Na + + e àna (h) Fe à Fe e (c) Ca e à Ca (f) S 2 às + 2e (i) Mn e àmn 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+ à Fe 3+ (f) Mn 3+ à Mn 4+ (b) K à K + (g) Cr 2+ à Cr 3+ (c) Sn 4+ à Sn 2+ (h) Cl 7+ à Cl 1+ (d) Cr 6+ à Cr 3+ (i) 3Cl2 à 6Cl (e) O2 à 2O (j) 4H + à 2H2 5. For each of the following equations, write the reduction half-reaction and the oxidation halfreaction. Reduction: (a) Cl2 + 2 KBr à 2 KCl + Br2 Oxidation: Reduction: (b) Cu + 2 Ag + à 2 Ag + Cu 2+ Oxidation: Reduction: (c) 2 Mg + O2 à 2 MgO Oxidation:

8 Practice: 6. Which half-reaction correctly represents oxidation? (1) Fe(s) à Fe 2+ (aq) + 2e (2) Fe 2+ (aq) à Fe(s) + 2e (3) Fe(s) + 2e à Fe 2+ (aq) (4) Fe 2+ (aq) + 2e à Fe(s) 7. Which equation shows a conservation of both mass and charge? 9. Given the balanced ionic equation: Which equation represents the oxidation halfreaction? (1) Zn(s) + 2e à Zn 2+ (aq) (2) Zn(s) à Zn 2+ (aq) + 2e (3) Cu 2+ (aq) à Cu(s) + 2e (4) Cu 2+ (aq) + 2e à Cu(s) (1) Cl2 + Br à Cl + Br2 (2) Cu + 2Ag + à Cu 2+ + Ag + (3) Zn + Cr 3+ à Zn 2+ + Cr (4) Ni + Pb 2+ à Ni 2+ + Pb 8. Which half-reaction equation represents the reduction of a potassium ion? (1) K + + e à K (3) K + à K + e (2) K + e à K + (4) K à K + + e 10. Given the equation: The reduction half-reaction is (1) Al à Al e (2) Cu e à Cu (3) Al + 3e à Al 3+ (4) Cu 2+ à Cu + 2e

9 Skill 4: Balancing Redox Equations Flipped Lesson o Redox reactions must be balanced by and. This is according to the Law of Conservation of matter and charge. Example: Balance the following reaction: Cu + HNO 3 Cu(NO 3 ) 2 + NO + H 2 O Step 1: Write the oxidation numbers for each atom in the reaction Cu+HNO 3 Cu(NO 3 ) 2 + NO+H 2 O Step 2: Identify the substance being oxidized and reduced. Oxidized: Reduced: Step 3: Write the half reactions: Step 4: Multiply each half reaction so the electron number is equal in both reactions. Step 5: Add the two balanced half-cell reactions together for NET Reaction. Practice: Balance the following redox reactions using the five-step procedure outlined above. 1. CsF + Na NaF + Cs 2. NaCl + Br2 NaBr + Cl2

10 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 + CuBr2 CrBr3 + Cu Write the half reactions and then balance the net equation! 1) Mg + Fe 3+ Mg 2+ + Fe 2+ Reduction: Oxidation: 2) Al + Ti 4+ Al 3+ + Ti Reduction: Oxidation: 3) Given the balanced equation: Mg(s) + Ni 2+ (aq) à 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 Practice: 1. Cu + Ag +1 à Cu +2 + Ag Reduction: Oxidation: 2. Fe + Pb +2 à Fe +3 + Pb Reduction: Oxidation: 3. Ag +1 Cr à Ag + Cr +3 Reduction: Oxidation: 4. Ni +2 + Li à Li +1 + Ni Reduction: Oxidation: 7. 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. (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.

12 Name Net reaction Bell Ringer Write half reactions: Combined for net reaction KCl + F2à KF + Cl2 Name Net reaction Bell Ringer Write half reactions: Combined for net reaction KCl + F2à KF + Cl2 Name Net reaction Bell Ringer Write half reactions: Combined for net reaction KCl + F2à KF + Cl2 Name Net reaction Bell Ringer Write half reactions: Combined for net reaction KCl + F2à KF + Cl2

13 NAME: Period Assign oxidation numbers to all the elements in each of the compounds or ions below: 1) KNO3 2) Mg3N2 3) H2SO4 Writing Half Reactions: 4. 2H2 + O2 à 2H2O Oxidized Reduced Balanced Half Reactions: Oxidation: Reduction: 5. Fe + Zn +2 à Fe +2 + Zn Oxidized Reduced Balanced Half Reactions: Oxidation: Reduction: 6. 2Ag +1 + Cu à Cu Ag Oxidized Reduced Balanced Half Reactions: Oxidation: Reduction:

14 Skill 5: Table J and Spontaneous Reactions o General Rule: elements on Table J are reactive than the elements below them. o Spontaneous rxn = rxn occurs w/out adding energy to system. If the single element is more active than the combined element, the reaction will be spontaneous. o Non-spontaneous rxn = rxn will not occur unless energy is added to system. If the single element is than the combined element, the reaction will NOT be spontaneous. Ex1: Zn + PbCl 2 à Ex2: Zn + BaO à For the following: State whether or not the reaction is spontaneous or non-spontaneous using Table J of the reference tables. 1) CsF + NaàNaF + Cs 2) MgCl 2 + CràMg + CrCl 3 3) Pb + Ag 3 PO 4 àpb 3 (PO 4 ) 2 + Ag 4. Which reaction occurs spontaneously? (1) Cl2(g) + 2NaBr(aq) à Br2(l) + 2NaCl(aq) (2) Cl2(g) + 2NaF(aq) à F2(g) + 2NaCl(aq) (3) I2(s) + 2NaBr(aq) à Br2(l) + 2NaI(aq) (4) I2(s) + 2NaF(aq) à F2(g) + 2NaI(aq) 5. Which metal reacts spontaneously with a solution containing zinc ions? (1) magnesium (3) copper (2) nickel (4) silver 6. Which metal with react with Zn 2+ spontaneously, but will not react with Mg 2+? (1) Al (3) Ni (2) Cu (4) Ba 7. Which of the following metals has the least tendency to undergo oxidation? (1) Ag (3) Zn (2) Pb (4) Li 8. Because tap water is slightly acidic, water pipes made of iron corrode over time, as shown by the balanced ionic equation below: 2Fe(s) + 6H + (aq) à 2Fe 3+ (aq) + 3H2(g) Explain, in terms of chemical reactivity, why copper pipes are less likely to corrode than iron pipes.

Skill 6: Electrochemical Batteries: Electrolytic and Voltaic Anode: The ACTIVE of the 2 metals (Table J) Spontaneously electrons to cathode The negative electrode in a VOLTAIC CELL Electrode where

VOLTAIC CELL Electrode where occurs (RED CAT) Electrolytic Cells: Cells that use ELECTRICAL ENERGY to force a chemical reaction to occur.

15 Skill 6: Electrochemical Batteries: Electrolytic and Voltaic Anode: The ACTIVE of the 2 metals (Table J) Spontaneously electrons to cathode The negative electrode in a VOLTAIC CELL Electrode where occurs (AN OX) Red Catà Reduction happens at cathode An Ox à Anode is where oxidation occurs Cathode: The ACTIVE of the 2 metals (Table J) Spontaneously electrons to it The positive electrode in a VOLTAIC CELL Electrode where occurs (RED CAT) Electrolytic Cells: Cells that use ELECTRICAL ENERGY to force a chemical reaction to occur. o The Process of Electroplating occurs through a forced reaction. o How do I know it s electrolytic? IF YOU SEE A POWER SOURCE CONNECTED. Use the diagram of a key being plated with copper to answer the following: 1. What is the name of the process shown in the diagram? 2. What is the purpose of the battery in this electrolytic cell? 3. Which electrode, A or B, attracts positive copper ions? 4. Given the reduction reaction for this cell: Cu 2+ (aq) + 2e à 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 5. 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

16 Voltaic Cells: Cells that convert CHEMICAL energy into energy. Number of containers Oxidation occurs at the: Reduction occurs at the: Anode charge Cathode charge Electrons flow from COMPARISON OF BATTERIES Voltaic Electrolytic 1. 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 Energy conversion that occurs in this cell Is this reaction spontaneous or does it require an outside power source to happen? 2. Which half-reaction can occur at the anode in a voltaic cell? (1) Ni e- Ni (2) Sn + 2e- Sn 2+ (3) Zn Zn e-

Skill 7: Voltaic Cell Problems, Flipped Lesson 1) Base your answers to the following questions on the diagram below, which represents a voltaic cell at 298 K and 1 atm.

17 Skill 7: Voltaic Cell Problems, Flipped Lesson 1) 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? (b) Write the balanced half-reaction equation that will occur in half-cell 1 when switch S is closed. (c) Describe the direction of electron flow between the electrodes when switch S is closed. Answer questions 2 and 3 using the diagram below, which represents an electrochemical cell. 2. When the switch is closed, in which half-cell does oxidation occur? 3. 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. 4) Aluminum is one of the most abundant metals in Earth s crust. The aluminum compound found in bauxite ore is Al2O3. 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, Na3AlF6, would dissolve bauxite ore. Electrolysis of the resulting mixture caused the aluminum ions in the Al2O3 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. (b) Write the balanced half-reaction equation for the reduction of Al 3+ to Al. (c) Explain, in terms of ions, why molten cryolite conducts electricity. (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.

18 5. 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. 6. 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. 7. Given the balanced equation representing a reaction occurring in an electrolytic cell: 2NaCl(l) à 2Na(l) + Cl2(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 occur 8. Base your answers to the following questions on the diagram of the voltaic cell below. (a) Identify the anode and the cathode. (b) Write the oxidation and reduction half-reactions for this voltaic cell. (f) State the electrode to which positive ions migrate when the switch is closed. (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.

48. Base your answers to the following questions on the information below. Underground iron pipes in contact with moist soil are likely to corrode.

The magnesium block acts as the anode and the iron pipe acts as the cathode. A diagram of this system is shown below.

19 48. 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. and zinc. (b) Explain, in terms of reactivity, why magnesium is preferred over zinc to protect underground iron pipes. Your response must include both magnesium 49. Base your answers to the following questions on the diagram of a voltaic cell and the balanced ionic equation below. (a) What is the total number of moles of electrons needed to completely reduce 6.0 moles of Ni 2+ (aq) ions? (b) Identify one metal from Reference Table J that is more easily oxidized than Mg(s). (c) Explain the function of the salt bridge in the voltaic cell.

21 Name Bell Ringer Identify one metal that does not react spontaneously with HCl(aq). Name Bell Ringer Identify one metal that does not react spontaneously with HCl(aq). Name Bell Ringer Identify one metal that does not react spontaneously with HCl(aq). Name Bell Ringer Identify one metal that does not react spontaneously with HCl(aq). Name Bell Ringer Identify one metal that does not react spontaneously with HCl(aq). Name Bell Ringer Identify one metal that does not react spontaneously with HCl(aq). Name Bell Ringer Identify one metal that does not react spontaneously with HCl(aq). Name Bell Ringer Identify one metal that does not react spontaneously with HCl(aq). Name Bell Ringer Identify one metal that does not react spontaneously with HCl(aq). Name Bell Ringer

22 Identify one metal that does not react spontaneously with HCl(aq). Name Bell Ringer Identify one metal that does not react spontaneously with HCl(aq). Name Bell Ringer Identify one metal that does not react spontaneously with HCl(aq).

Practice Packet Unit 12: Electrochemistry

Regents Chemistry: PRACTICE PACKET: ELECTROCHEMISTRY Practice Packet Unit 12: Electrochemistry Redox and Batteries? Ain t nobody got time for that!!! 1 For each word, provide a short but specific definition