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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) H = +1 H = -1

  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 Na = +1O = -2 Therefore, N = +5

  12. Examples: Assign oxidation #’s to each element: b) SO32- +4 -2 O = -2, therefore S must = +4 to balance the charges and have an overall charge of 2-

  13. Examples: Assign oxidation #’s to each element: c) HCO3- +1 +4 -2

  14. Examples: Assign oxidation #’s to each element:Do on your own: • H3PO4 • Cr2O72- • K2Sn(OH)6

  15. 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)

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

  17. Are all rxns REDOX rxns?You must determine this… • a reaction is “redox” if a change in oxidation # happens; if no change in oxidation # occurs, the reaction is nonredox.

  18. Examples MgCO3 MgO + CO2 MgCO3 is an ionic compound, so what is Mg’s charge in an ionic compound? +2 -2 +4 -2 +4 +2 -2 The carbonate ion CO32- is the other ion, let’s figure out C because we already know O. Is this a redox or nonredox reaction? NONREDOX (no change in oxidation numbers)

  19. Examples Zn + CuSO4 ZnSO4 + Cu Which oxidation numbers do we already know? +6 +2 -2 +2 +6 -2 0, free element 0 Break down this ionic compound into its ions Is this a redox or nonredox reaction? Cu and SO42- So, Cu must be Cu2+ REDOX reaction O = -2 in SO42-, so S must be…?

  20. Examples NaCl + AgNO3 AgCl + NaNO3 Redox or nonredox?

  21. Examples CO2 + H2O  C6H12O6 + O2 +1 -2 +1 +4 -2 0 -2 0 What happened to Carbon? It went from +4 oxidation # to 0. Was Carbon oxidized or reduced? REDUCED OIL RIG (oxidation is losing electrons so oxidation number increases, where as reduction is gaining electrons so oxidation number decreases)

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