(5) Balance each Redox equation using the half reaction method. Identify the oxidizing agent and the reducing agent. (a) Zn + Ag + Zn 2+ + Ag

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1 Oxidation Numbers and Redox Reactions (1) Assign oxidation numbers to each element in the following compounds/ions. Name: Period: (a) HNO 3 (i) CO 2 + (b) NH 4 + (j) N 2 H 5 2- (c) CrO 4 (k) H 4 P 2 O 7 (d) HClO 3 2- (l) SO 4 (e) S 2 F 10 (m) LiH 2- (f) C 2 O 4 (n) Na 2 O 2 (g) FeO (o) C 3 H 8 - (h) IO 4 (p) K 2 PtCl 6 (2) Assign oxidation numbers to each element. Determine if the reaction is a redox reaction. (a) N 2 + 3H 2 2NH 3 (b) CaBr 2 + 2AgNO 3 2AgBr + Ca(NO 3 ) 2 (c) 3CuSO 4 + 2Al Al 2 (SO 4 ) 3 + 3Cu (d) FeCl 2 + 2NaOH Fe(OH) 2 + 2NaCl (e) CH 4 + 2O 2 CO 2 + 2H 2 O (3) Determine if each change represents oxidation or reduction. (a) K K + (b) Br 2 Br - (c) Cr 2+ Cr 3+ (d) U 6+ U 4+ (4) Assign oxidation numbers to each atom and determine which element is being oxidized and which element is being reduced. Identify the oxidizing agent and the reducing agent. (a) Cr 2 O HNO 2 + 5H + 2Cr NO H 2 O (b) 2IO N 2 O + 2H + I NO + H 2 O (c) 4MnO Te + 2OH - 4MnO 2 + 3TeO H 2 O (d) 3ClO I - 3Cl - + 4IO 3 - (e) 3IO PH 3 3I - + 2P H 2 O (f) P 4 + 2NO OH - + 3H 2 O 4H 2 PO N 2 O (5) Balance each Redox equation using the half reaction method. Identify the oxidizing agent and the reducing agent. (a) Zn + Ag + Zn 2+ + Ag (d) Al + Fe 2+ Al 3+ + Fe (b) F 2 + Na F - + Na + (e) I 2 + Cr I - + Cr 3+ (c) Cr 2+ + Sn 4+ Cr 3+ + Sn 2+

2 (6) Balance each Redox equation occurring in Acidic Solution. Identify the oxidizing agent and the reducing agent. (a) Cl - + NO Cl 2 + N 2 O (d) BrO Mn 2+ Br 2 + MnO 2 (b) PbO 2 + I 2 Pb IO 3 (e) P 4 + IO - 3 H 2 PO I - (c) Cr 2 O S Cr 3+ + H 2 SO 3 (f) MnO C 2 O 4 2- Mn 2+ + CO 2 (7) Balance each Redox equation occurring in Basic Solution. Identify the oxidizing agent and the reducing agent (a) MnO SO MnO 2 + SO 4 (d) NO AsO 3- AsO NO (b) IO Cl - Cl 2 + I 2 (e) CrO ClO - 4 CrO Cl - (c) Cr + ClO - 4 CrO ClO 3 (f) TeO N 2 O 4 Te + NO 3 (8) Balance the following Redox equation occurring in acidic solution: Br 2 BrO Br - (9) Balance the following Redox equation occurring in basic solution: N 2 O N 2 H 4 + NO 3 - Answers: (1) (a) H: +1 N: +5 O: -2 (b) N: -3 H: +1 (c) Cr: 6+ O: 2- (d) H: +1 Cl: 5+ O: -2 (e) S: +5 F: -1 (f) C: 3+ O: -2 (g) Fe: +2 O: -2 (h) I: +7 O: -2 (i) C:+4 O: -2 (j) N: -2 H: +1 (k) H: +1 P: +5 O: -2 (l) S: +6 O: -2 (m) Li: +1 H: -1 (n) Na: +1 O: -1 (o) C: -8/3 H: +1 (p) K: +1 Pt: +4 Cl: -1 (2) (a) redox (b) not redox (c) redox (d) not redox (e) redox (3) (a) oxidation (b) reduction (c) oxidation (d) reduction (4) (a) ox agent: Cr 2 O 7 2- red agent: HNO 2 (b) ox agent: IO 3 - red agent: N 2 O (c) ox agent: MnO 4 - red agent: Te (d) ox agent: ClO 4 - red agent: I - (e) ox agent: IO 4 - red agent: PH 3 (f) ox agent: NO 2 red agent: P 4 (5) (a) Zn + 2Ag + Zn Ag (ox agent: Ag + red agent: Zn) (b) F 2 + 2Na 2F - + 2Na + (ox agent: F 2 red agent: Na) (c) 2Cr 2+ + Sn 4+ 2Cr 3+ + Sn 2+ (ox agent: Sn 4+ red agent: Cr 2+ ) (6) (a) 2Cl - + 2NO + 2H + Cl 2 + N 2 O + H 2 O (ox agent: NO red agent: Cl - ) (b) 5PbO 2 + I 2 + 8H + 5Pb IO H 2 O (ox agent: PbO 2 red agent: I 2 ) (c) 2Cr 2 O S + 16H + 4Cr H 2 SO 3 + 5H 2 O (ox agent: Cr 2 O 7 2- red agent: S) (7) (a) 2MnO SO H 2 O 2MnO 2 + 3SO OH - (ox agent: MnO 4 - red agent: SO 3 2- ) (b) 2IO Cl - + 6H 2 O 5Cl 2 + I OH - (ox agent: IO 3 - red agent: Cl - ) (c) 2Cr + 3ClO OH - 2CrO ClO H 2 O (ox agent: ClO 4 - red agent: Cr) (d) 2Al + 3Fe 2+ 2Al Fe (ox agent: Fe 2+ red agent: Al) (e)3 I 2 + 2Cr 6I - + 2Cr 3+ (ox agent: I 2 red agent: Cr) (d) 2BrO Mn H 2 O Br 2 + 5MnO 2 + 8H + (ox agent: BrO 3 - red agent: Mn 2+ ) (e) 3P IO H 2 O 12H 2 PO I H + (ox agent: IO 3 - red agent: P 4 ) (f) 2MnO C 2 O H + 2Mn CO 2 + 8H 2 O (ox agent: MnO 4 - red agent: C 2 O 4 2- ) (d) 2NO AsO 3- + H 2 O AsO NO + 2OH - (ox agent: NO 3 - red agent: AsO 3- ) (e) 8CrO ClO OH - 8CrO Cl - + 4H 2 O (ox agent: ClO 4 - red agent: CrO 2 - ) (f) TeO N 2 O 4 + 2OH - Te + 4NO H 2 O (ox agent: TeO 3 2- red agent: N 2 O 4 ) (8) 3Br 2 + 3H 2 O BrO Br - + 6H + (9) 7N 2 O + 5H 2 O + 6OH - 4N 2 H 4 + 6NO 3 -

3 Redox Titrations Name: Period: (1) In acidic solution, MnO 4 - reacts with Sn 2+ to produce Mn 2+ and Sn 4+ according to the following balanced equation: 16H + + 5Sn 2+ +2MnO 4-5Sn Mn H 2 O What volume of M KMnO 4 solution is need to react with 35.0 ml of M SnCl 2 solution? (2) SO 3 2- reacts with ClO 3 - to produce SO 4 2- and Cl - according to the following balanced equation: 3SO ClO 3-3SO Cl - What is the concentration if 50.0 ml of Na 2 SO 3 solution requires 25.0 ml of M NaClO 3 to titrate? (3) A sample of iron ore weighing 2.00 g was dissolved in acid to produce Fe 2+ ions. The solution was then titrated with ml of M KMnO 4 solution. In acidic solution, Fe 2+ reacts with MnO 4 - to produce Fe 3+ and Mn 2+. (a) Write the balanced chemical equation for this reaction. (b) Determine the mass of pure iron in the sample of iron ore. (c) Determine the percent iron in the sample of iron ore. (4) A 1.00 g sample of hydrogen peroxide (H 2 O 2 ) solution is titrated with ml of M KMnO 4 solution. In acidic solution, H 2 O 2 reacts with MnO 4 - to produce O 2 and Mn 2+. (a) Write the balanced chemical equation for this reaction. (b) Determine the mass of pure hydrogen peroxide in the solution of hydrogen peroxide. (c) Determine the percent hydrogen peroxide in the solution. (5) A 1.00 g sample of solid containing HNO 2 is titrated with 11.6 ml of M KMnO 4 solution. In acidic solution, HNO 2 reacts with MnO 4 - to produce NO 3 - and Mn 2+. (a) Write the balanced chemical equation for this reaction. (b) Determine the mass of pure nitrous acid in the sample. (c) Determine the percent nitrous acid in the sample. Answers: (1) 28.0 ml (2) M (3) (a) 5Fe 3+ + MnO H + 5Fe 2+ + Mn H 2 O (b) g Fe (c) 38.4% (4) (a) 5H 2 O 2 + 2MnO H + 5O 2 + 2Mn H 2 O (b) g H 2 O 2 (c) 29.9% (5) (a) 5HNO 2 + 2MnO H + 5NO Mn H 2 O (b) g HNO 2 (c) 6.82%

4 Predicting Redox Reactions Complete and Balance the following reactions Name: Period: (1) Copper (II) sulphate is combined with aluminum metal. (2) A solution of iron (II) chloride is mixed with acidified sodium dichromate solution. (3) Cadmium metal is placed in a solution of tin (II) chloride (4) A solution of tin (II) nitrate is added to a solution of silver nitrate. (5) Chlorine gas is added to sodium iodide solution. (6) Manganese (IV) oxide is added to sodium bromide. (7) Zinc metal is placed in a solution of nickel (II) nitrate (8) A strip of aluminum metal is placed in liquid bromine. (9) A bar of strontium metal is placed in copper (II) nitrate solution (10) Tin (II) nitrate solution is mixed with acidified potassium permanganate solution. Answers: (1) 3Cu Al 3Cu + 2Al 3+ (2) 6Fe 2+ + Cr 2 O H + 6Fe Cr H 2 O (3) Cd + Sn 2+ Cd 2+ + Sn (4) Sn Ag + Sn Ag (5) Cl 2 + 2I - 2Cl - + I 2 (6) MnO 2 + 2Br - + 4H + Mn 2+ + Br 2 + 2H 2 O (7) Zn + Ni 2+ Zn 2+ + Ni (8) 2Al + 3Br 2 2Al Br - (9) Sr + Cu 2+ Sr 2+ + Cu (10) 5Sn MnO H + 5Sn Mn H 2 O

5 Electrochemical Cells Name: Period: (1) Consider the following electrochemical cell. (a) Write the cathode half reaction. Label the cathode. (b) Write the anode half reaction. Label the anode. (d) Calculate the cell voltage. (Eº cell ) (f) Which electrode loses mass and which electrode gains mass? (g) Give the line notation for this electrochemical cell. (2) Consider the following electrochemical cell. (a) Write the cathode half reaction. Label the cathode. (b) Write the anode half reaction. Label the anode. (d) Calculate the cell voltage. (Eº cell ) (f) Which electrode loses mass and which electrode gains mass? (g) Give the line notation for this electrochemical cell. (3) Consider the following electrochemical cell. (a) Write the cathode half reaction. Label the cathode. (b) Write the anode half reaction. Label the anode. (d) Calculate the cell voltage. (Eº cell ) (f) Which electrode loses mass and which electrode gains mass? (g) Give the line notation for this electrochemical cell.

6 (4) Consider the following electrochemical cell. (a) Write the cathode half reaction. Label the cathode. (b) Write the anode half reaction. Label the anode. (d) Calculate the cell voltage. (Eº cell ) (f) Give the line notation for this electrochemical cell. (5) Consider the following electrochemical cell. (a) Write the cathode half reaction. Label the cathode. (b) Write the anode half reaction. Label the anode. (d) Calculate the cell voltage. (Eº cell ) (f) Give the line notation for this electrochemical cell. (6) Consider the following electrochemical cell. (a) Write the cathode half reaction. Label the cathode. (b) Write the anode half reaction. Label the anode. (d) Calculate the cell voltage. (Eº cell ) (f) Give the line notation for this electrochemical cell.

7 (7) Consider the following electrochemical cell. The cell voltage is 0.47 V. (a) What metal is the missing electrode composed of? What is a possible solution for that half-cell? (b) Write the cathode half reaction. Label the cathode. (c) Write the anode half reaction. Label the anode. (d) Write the overall reaction. (f) Which electrode loses mass and which electrode gains mass? (g) Give the line notation for this electrochemical cell. (8) Consider the following electrochemical cell. The cell voltage is 2.07 V. (a) What metal is the missing electrode composed of? What is a possible solution for that half-cell? (b) Write the cathode half reaction. Label the cathode. (c) Write the anode half reaction. Label the anode. (d) Write the overall reaction. (f) Give the line notation for this electrochemical cell.

8 Answers: (1) (a) Sn e - Sn (b) Zn Zn e - (c) Sn 2+ + Zn Sn + Zn 2+ (d) 0.62 V (e) e - NO 3- K + (f) anode (Zn) loses mass, cathode (Sn) gains mass (g) Zn (s) Zn 2+ (aq) Sn 2+ (aq) Sn (s) (2) (a) Cu e - Cu (b) Al Al e - (c) 3Cu Al 3 Cu + 2Al 3+ (d) 2.00 V (e) e - NO 3- K + (f) anode (Al) loses mass, cathode (Cu) gains mass (g) Al (s) Al 3+ (aq) Cu 2+ (aq) Cu (s) (3) (a) Ag + + e - Ag (b) Ni Ni e - (c) 2Ag + + Ni 2Ag + Ni 2+ (d) 1.05 V (e) e - NO 3- K + (f) anode (Ni) loses mass, cathode (Ag) gains mass (g) Ni (s) Ni 2+ (aq) Ag + (aq) Ag (s) (4) (a) Ag + + e - Ag (b) Cr 2+ Cr 3+ + e - (c) Ag + + Cr 2+ Ag + Cr 3+ (d) 1.21 V (e) e - NO 3- K + (f) Pt (s) Cr 3+ (aq), Cr 2+ (aq) Ag + (aq) Ag (s) (5) (a) Cl 2 + 2e - 2Cl - (b) Pb Pb e - (c) Cl 2 + Pb 2Cl - + Pb 2+ (d) 1.49 V (e) e - NO 3- K + (f) Pb (s) Pb 2+ (aq) Cl - (aq) Cl 2 (g) Pt (s) (6) (a) Hg e - Hg (b) Cd Cd e - (c) Hg 2+ + Cd Hg + Cd 2+ (d) 1.25 V (e) e - NO 3- K + (f) Cd 2+ (aq) Cd (s) Hg 2+ (aq) Pt (s) Hg (l) (7) (a) Cu, Cu(NO 3 ) 2 (b) Cu e - Cu (c) Pb Pb e - (d) Cu 2+ + Pb Cu + Pb 2+ (e) e - NO 3- K + (f) anode (Pb) loses mass, cathode (Cu) gains mass (g) Pb (s) Pb 2+ (aq) Cu 2+ (aq) Cu (s) (8) (a) Ni, Ni(NO 3 ) 2 (b) Co 3+ + e - Co 2+ (c) Ni Ni e - (d) 2Co 3+ + Ni 2Co 2+ + Ni 2+ (e) e - NO 3- K + (f) Ni 2+ (aq) Ni (s) Co 2+ (aq), Co 3+ (aq) Pt (s)

9 Electrolytic Cells Name: Period: (1) Consider the following electrolytic cell. (a) Write the cathode half reaction. (b) Write the anode half reaction. (d) Calculate the power that must be supplied to operate the cell. (2) Consider the following electrolytic cell. (a) Write the cathode half reaction. (b) Write the anode half reaction. (d) Calculate the power that must be supplied to operate the cell. (3) Consider the following electrolytic cell. (a) Write the cathode half reaction. (b) Write the anode half reaction. (d) Calculate the power that must be supplied to operate the cell.

10 (4) Consider the following electrolytic cell. (a) Write the cathode half reaction. (b) Write the anode half reaction. (d) Calculate the power that must be supplied to operate the cell. (5) Consider the following electrolytic cell. (a) Write the cathode half reaction. (b) Write the anode half reaction. (d) Calculate the power that must be supplied to operate the cell. (6) Consider the following electrolytic cell. (a) Write the cathode half reaction. (b) Write the anode half reaction. (d) Calculate the power that must be supplied to operate the cell.

11 (7) Consider the half reaction: Al e - Al. What mass of Aluminum is produced if 6.00 A is applied over 20.0 minutes? (8) Consider the half reaction: Cu e - Cu. What current is required to produce 0.50 g of Cu in 15.0 minutes? (9) Consider the half reaction: Pb e - Pb. How many minutes must a 10.0 A current be applied to produce 4.50 g of lead metal?

12 Answers: (1) (a) Na + + e - Na (b) 2F - F 2 + 2e - (c) 2Na + + 2F - 2Na + F 2 (d) 5.58 V (2) (a) Zn e - Zn (b) 2I - I 2 + 2e - (c) Zn F - Zn + I 2 (d) 1.29 V (3) (a) Ni e - Ni (b) Ag Ag + + e - (c) Ni Ag Ni + 2Ag + (d) 1.05 V (4) (a) K + + e - K (b) 2Cl - Cl 2 + 2e - (c) 2K + + 2Cl - 2K + Cl 2 (d) 4.28 V (5) (a) 2H 2 O + 2e - H 2 + 2OH - (b) 2Br - Br 2 + 2e - (c) 2H 2 O + 2Br - H 2 + 2OH - + Br 2 (d) 1.90 V (6) (a) 2H 2 O + 2e - H 2 + 2OH - (b) Ni Ni e - (c) 2H 2 O + Ni H 2 + 2OH - + Ni 2+ (d) 0.58 (7) g (8) 1.7 A (9) 6.99 minutes

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