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UNIT V

UNIT V. Balancing Redox Reactions and Redox Titrations. Preparing a Redox Table. Example: Four metals A, B, C, & D were tested with separate solutions of A 2+ , B 2+ , C 2+ & D 2+ . Some of the results are summarized in the following table:

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UNIT V

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  1. UNIT V Balancing Redox Reactions and Redox Titrations

  2. Preparing a Redox Table Example: Four metals A, B, C, & D were tested with separate solutions of A2+, B2+, C2+ & D2+. Some of the results are summarized in the following table: List the ions in order from the strongest to weakest oxidizing agent.

  3. Preparing a Redox Table NOTE: For the same element: the more positive species is always the Oxidizing Agent. • Ex. A2+ vs. A Using data: • Since B2+ does not oxidize A, B2+ must be below A on the table. • Since C2+ reacts with A, C2+ must be above A. •  Since A2+ reacts with D, A2+ must be above D on the table. But is D2+ above or below B2+? We don’t know yet. • D2+ does not react with B. Now we know that D2+ must be below B on the table.

  4. Preparing a Redox Table • So now we have our complete table: • So now we have our answer; The ions in order of strongest to weakest oxidizing agent is: C2+, A2+, B2+, D2+ • The order of reducing agent from strongest to weakest is: D, B, A, C.

  5. Preparing a Redox Table Example: Four non-metal oxidizing agents X2, Y2, Z2 and W2 are combined with solutions of ions: X-, Y-, Z- and W-.   The following results were obtained; • X2 reacts with W- and Y- only. • Y- will reduce W2 List the reducing agents from strongest to weakest

  6. Preparing a Redox Table • X2 will be above W- & Y-, but below Z- •  Since Y- reduces W2, Y- must be lower on the right of W2. • The reducing agents from strongest to weakest are: Y‑, W‑, X‑, Z‑

  7. Preparing a Redox Table Example: Four solutions A(NO3)2, B(NO3)2, C(NO3)2, and D(NO3)2 are added to metals, A, B, C, & D. The following information is found: • The metal A will not react with any of the solutions • C(NO3)2 reacts spontaneously with B • B will not react with D(NO3)2 • Make a small reduction table showing reductions of the metallic ions. (Don’t forget to discard the spectator nitrate ions.)

  8. Preparing a Redox Table • List the oxidizing agents in order of strongest to weakest: _________________________ • List the reducing agents in order of strongest to weakest: _________________________ • Would it be safe to store A(NO3)2 solution in a container made of the metal D? ______________ • Hebden Textbook p. 200 Questions #14-18

  9. Balancing Half-Reactions • Some half-rxns are on the table, but not all! • Given the solution: Is soln acidic or basic? Think: Major Hydroxide (Major  O  H  - charge) (Major atoms atoms other than O & H)

  10. Balancing Half-Reactions Ex. S2O82- HSO3- (acid soln) 1. Balance Major Atoms (S in this case) S2O82- 2HSO3- 2. Balance “O” atoms by adding H2O (to the side with less O’s)  S2O82- 2HSO3- + 2H2O 3. Balance “H” atoms by adding H+ (to the side with less H’s) S2O82- + 6H+ 2HSO3- + 2H2O 4. Balance charge by adding e-‘s (to the more + side)  S2O82- + 6H+ 2HSO3- + 2H2O TIC = +4 TIC = -2

  11. Balancing Half-Reactions So the final balanced half-rxn is: S2O82- + 6H+ + 6 e- 2HSO3- + 2H2O TIC = -2 TIC = -2

  12. Balancing Half-Reactions Example: MnO4- Mn2+ (acid soln)

  13. Balancing Half-Reactions In Basic Solution: •  Do the first steps of the balancing just like an acid. • Ex. MnO2 MnO4- (basic solution) • Major (Mn already balanced) • Oxygen 2H2O + MnO2 MnO4- • Hydrogen 2H2O + MnO2 MnO4- + 4H+ • Charge 2H2O + MnO2 MnO4- + 4H+ + 3e-

  14. Balancing Half-Reactions In basic solution, write the reaction: H+ + OH- H2O or H2O  H+ + OH- in whichever way is needed to cancel out the H+’s and add to the half-rxn. • You must write the whole water equation • We need 4H+ on the left to cancel the 4H+ on the right side 2H2O + MnO2 MnO4- + 4H+ + 3e- 4H+ + 4OH- 4H2O   ________________________________________ MnO2 + 4OH- MnO4- + 2H2O + 3e-

  15. Balancing Half-Reactions Example: Pb  HPbO2- (basic soln)

  16. Balancing Overall Redox Reactions Half-Reaction Method • Break up rxn into 2 half-rxns. • Balance each one (in acidic or basic as given). • Multiply each half rxn by whatever is needed to cancel out e-‘s. (Note: balanced half-rxns show e-‘s (on left for reduction - on right for oxidation) but balanced redoxrxns don’t show e-‘s). • Add the 2 half-rxns canceling e-‘s and anything else (usually H2O’s, H+’s or OH-‘s) in order to simplify.

  17. Balancing Overall Redox Reactions Half-Reaction Method Example: U4+ + MnO4- Mn2+ + UO22+ (acidic)

  18. Balancing Overall Redox Reactions Half-Reaction Method To simplify: • Take away 10e- from both sides • Take away 16H+’s from both sides • Take away 8H2O’s from both sides 5U4+ + 2H2O + 2MnO4- 5UO22+ + 4H+ + 2Mn2+ • Quick check by finding TIC’s on both sides. • If you have time check all atoms also. If TIC’s are not equal, you messed up!

  19. Balancing Overall Redox Reactions Half-Reaction Method Example: SO2 + IO3- SO42- + I2 (basic soln)

  20. Balancing Overall Redox Reactions Half-Reaction Method Quick Notes •  Some redox equations have just one reactant. • Use this as the reactant in both half-rxns. • These are called “Self-Oxidation-Reduction” or “Disproportionation Reactions.” Ex. Br2 Br-- + BrO3- (basic soln)

  21. Balancing Overall Redox Reactions Half-Reaction Method • Half-rxns are: Br2 Br- Br2 BrO3- • Hebden Textbook p. 207 Question #24

  22. Balancing Overall Redox Reactions Oxidation Number Method • This is optional • As long as one method (not guessing!) works for you, that’s fine • Read examples on page 208-209 in Hebden Textbook • Review Question #25

  23. Redox Titrations • Same as in other units (solubility/acids-bases) • Coefficient ratios for the “mole bridge” are obtained by the balanced redox equation. MOLE BRIDGE

  24. Redox Titrations Example: Acidified hydrogen peroxide (H2O2) is used to titrate a solution of MnO4- ions of unknown concentration. Two products are O2 gas and Mn2+. • Write the balanced redox equation:

  25. Redox Titrations b) It takes 6.50 mL of 0.200 M H2O2 to titrate a 25.0 mL sample of MnO4- solution. Calculate the original [MnO4-].

  26. Finding a Suitable Standard • Use redox table: • If sample is on the left (OA)…use something below it on the right (RA) • If sample is on the right (RA)…use something above it on the left (OA) • Good standards will change colour as they react • Acidified MnO4- (purple) = Mn2+ (clear) • Acidified Cr2O72- (orange) = Cr3+ (pale green) http://jchemed.chem.wisc.edu/JCESoft/CCA/CCA3/MVHTM/TITREDO/TITR1.HTM http://jchemed.chem.wisc.edu/JCESoft/CCA/CCA3/MVHTM/TITREDO/TITR2.HTM • Hebden Textbook p. 213-214 Questions #26, 29

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