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Redox Reactions

Redox Reactions. Chapter 18. + O 2 . Oxidation-Reduction (Redox) Reactions. “redox” reactions: rxns in which electrons are transferred from one species to another oxidation & reduction always occur simultaneously we use OXIDATION NUMBERS to keep track of electron transfers.

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Redox Reactions

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  1. Redox Reactions Chapter 18 + O2

  2. Oxidation-Reduction (Redox) Reactions • “redox” reactions: rxns in which electrons are transferred from one species to another • oxidation & reduction always occur simultaneously • we use OXIDATION NUMBERS to keep track of electron transfers

  3. Rules for Assigning Oxidation Numbers: 1) the ox. state of any free (uncombined) element is zero. • Ex: Na, S, O2, H2, Cl2, O3

  4. Rules for Assigning Oxidation Numbers: 2) The ox. state of an element in a simple ion is the charge of the ion. Mg2+ oxidation of Mg is +2

  5. Rules for Assigning Oxidation Numbers: • 3) the ox. # for hydrogen is +1 (unless combined with a metal, then it has an ox. # of –1) Ex: NaOH (H bonded to O) v. NaH (H bonded to Na)

  6. Rules for Assigning Oxidation Numbers: 4) the ox. # of fluorine is always –1.

  7. Rules for Assigning Oxidation Numbers: 5) the ox. # of oxygen is usually –2. Why USUALLY? Not -2 when it’s in a peroxide, such as hydrogen peroxide: H2O2

  8. Rules for Assigning Oxidation Numbers: 6) in any neutral compound, the sum of the oxidation #’s = zero.

  9. Rules for Assigning Oxidation Numbers: 7) in a polyatomic ion, the sum of the oxidation #’s = the overall charge of the ion.

  10. Rules for Assigning Oxidation Numbers: **use these rules to assign oxidation #’s; assign known #’s first, then fill in the #’s for the remaining elements:

  11. Examples: Assign oxidation #’s to each element: a) NaNO3

  12. Examples: Assign oxidation #’s to each element: b) SO32-

  13. Examples: Assign oxidation #’s to each element: c) HCO3-

  14. Examples: Assign oxidation #’s to each element: d) H3PO4

  15. Examples: Assign oxidation #’s to each element: e) Cr2O72-

  16. Examples: Assign oxidation #’s to each element: f) K2Sn(OH)6

  17. Definitions • Oxidation: the process of losing electrons (ox # increases) • Reduction: the process of gaining electrons (ox # decreases) • Oxidizing agents: species that cause oxidation (they are reduced) • Reducing agents: species that cause reduction (they are oxidized)

  18. To help you remember… OIL RIG • Oxidation Is Loss • Reduction Is Gain

  19. Are all rxns REDOX rxns? • a reaction is “redox” if a change in oxidation # happens; if no change in oxidation # occurs, the reaction is nonredox.

  20. Examples MgCO3 MgO + CO2

  21. Examples Zn + CuSO4 ZnSO4 + Cu

  22. Examples NaCl + AgNO3 AgCl + NaNO3

  23. Examples CO2 + H2O  C6H12O6 + O2

  24. Balancing Redox Equations

  25. Balancing Redox Equations • In balancing redox equations, the # of electrons lost in oxidation (the increase in ox. #) must equal the # of electrons gained in reduction (the decrease in ox. #) • There are 2 methods for balancing redox equations: • Change in Oxidation-Number Method • The Half-Reaction Method

  26. 1. Change in Oxidation-Number Method: • based on equal total increases and decreases in oxidation #’s Steps: • Write equation and assign oxidation #’s. • Determine which element is oxidized and which is reduced, and determine the change in oxidation # for each. • Connect the atoms that change ox. #’s using a bracket; write the change in ox. # at the midpoint of each bracket. • Choose coefficients that make the total increase in ox. # = the total decrease in ox. #. • Balance the remaining elements by inspection.

  27. S + HNO3 SO2 + NO + H2O Example

  28. If needed, reactions that take place in acidic or basic solutions can be balanced as follows:

  29. Example: Balance the following equation, assuming it takes place in acidic solution. ClO4- + I- Cl- + I2

  30. Example: Balance the following equation, assuming it takes place in basic solution. ClO4- + I- Cl- + I2

  31. 2. The Half-Reaction Method: • separate and balance the oxidation and reduction half-reactions. Steps: • Write equation and assign oxidation #’s. • Determine which element is oxidized and which is reduced, and determine the change in oxidation # for each. • Construct unbalanced oxidation and reduction half reactions. • Balance the elements and the charges (by adding electrons as reactants or products) in each half-reaction. • Balance the electron transfer by multiplying the balanced half-reaction by appropriate integers. • Add the resulting half-reaction and eliminate any common terms to obtain the balanced equation.

  32. Example: Balance the following using the half-reaction method: HNO3 + H2S  NO + S + H2O

  33. If needed, reactions that take place in acidic solutions can be balanced as follows:

  34. If needed, reactions that take place in basic solutions can be balanced as follows:

  35. HOMEWORK: Balance the following using the half-rxn method… In acidic sol’n: a) Cu + NO3- Cu2+ + NO b) Cr2O72- + Cl-  Cr3+ + Cl2 c) Pb + PbO2 + H2SO4 PbSO4 In basic sol’n: a) Al + MnO4- MnO2 + Al(OH)4- b) Cl2 Cl- + OCl- c) NO2- + Al  NH3 + AlO2-

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