1 / 35

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.

kitty
Télécharger la présentation

Redox Reactions

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  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-

More Related