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Oxidation-Reduction Reactions

Oxidation-Reduction Reactions. (electron transfer reactions). (redox reactions). 2Mg ( s ) + O 2 ( g ) 2MgO ( s ). 2Mg 2Mg 2+ + 4e -. O 2 + 4e - 2O 2-. 2Mg + O 2 + 4e - 2Mg 2+ + 2O 2- + 4e -. 2Mg + O 2 2MgO.

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Oxidation-Reduction Reactions

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  1. Oxidation-Reduction Reactions (electron transfer reactions) (redox reactions)

  2. 2Mg (s) + O2 (g) 2MgO (s) 2Mg 2Mg2+ + 4e- O2 + 4e- 2O2- 2Mg + O2 + 4e- 2Mg2+ + 2O2- + 4e- 2Mg + O2 2MgO Oxidation-Reduction Reactions (electron transfer reactions) (redox reactions) Oxidation half-reaction (lose e-) Reduction half-reaction (gain e-)

  3. 2Mg (s) + O2 (g) 2MgO (s) 2Mg 2Mg2+ + 4e- O2 + 4e- 2O2- Oxidation-Reduction Reactions (electron transfer reactions) Oxidation half-reaction (lose e-) Reduction half-reaction (gain e-) L E O G E R

  4. 2Mg (s) + O2 (g) 2MgO (s) 2Mg 2Mg2+ + 4e- O2 + 4e- 2O2- Oxidation-Reduction Reactions (electron transfer reactions) Oxidation half-reaction (lose e-) Reduction half-reaction (gain e-) ? ? ? ? Reducing agent causes oxidation Oxidizing agent causes reduction

  5. Oxidation number The charge the atom would have in a molecule (or an ionic compound) if electrons were completely transferred. • Free elements (uncombined state) have an oxidation number of zero. Na, Be, K, Pb, H2, O2, P4 = 0 • In monatomic ions, the oxidation number is equal to the charge on the ion. Li+ = +1; Fe3+ = +3; O2- = -2 • The oxidation number of oxygen isusually–2. In H2O2 and O22- it is –1. Don’t forget the free element rule for O2.

  6. The oxidation number of hydrogen is +1except when it is bonded to metals in binary compounds (hydride). In these cases, its oxidation number is –1. • Fluorine is always –1. (except elemental F2), Other halogens have negative numbers when occurring as halides (eg Cl-). In oxoacids and oxoanions, they have positive numbers. 6. The sum of the oxidation numbers of all the atoms in a molecule or ion is equal to the charge on the molecule or ion. 7. Oxidation numbers do not have to be integers (but they usually are in this class). One common example is the oxidation number of O = -1/2 in superoxide, O2-.

  7. Figure 4.10 The oxidation numbers of elements in their compounds

  8. 2Mg (s) + O2 (g) 2MgO (s) Oxidation Number 0 0 +2 -2 • Rules: • Elements = zero • Sum of oxidation numbers = charge on molecule • Sum of oxidation numbers reactants = sum of oxidation numbers products

  9. Harder examples: IF7 F = -1 ? ? ? 7x(-1) + ? = 0 I = +7 K2Cr2O7 NaIO3 O = -2 K = +1 O = -2 Na = +1 3x(-2) + 1 + ? = 0 7x(-2) + 2x(+1) + 2x(?) = 0 I = +5 Cr = +6

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