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Redox reactions

Redox reactions

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Redox reactions

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  1. Redox reactions Volumetric Analysis Unit 3

  2. Types of chemical reactions Fall into several categories. 1) Precipitation reactions The general form of the equation for such a reaction is: AB + CD  AD + CB 2) Acid-Base reactions. An acid substance react with a substance called a base. Such reactions involve the transfer of a proton (H+) between reactants. 3) Oxidation-reduction reactions. These involve the transfer of electrons between reactants.

  3. Oxidation: Reduction: Gain of oxygen Loss of oxygen Loss of electrons Gain of electrons Increase in oxidation number Decrease in oxidation number

  4. You can’t have one without the other! • Reduction (gaining electrons) can’t happen without an oxidation to provide the electrons. • You can’t have 2 oxidations or 2 reductions in the same equation. Reduction has to occur at the cost of oxidation • Oxidation and reductionalways occur simultaneously

  5. 2Mg(s) + O2(g)  2MgO(s) Mg  Mg2+ +2e- Oxidised – loss of e- O2 O2- Reduced – gain of e- +2e-

  6. Review of Oxidation numbers 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+, Li = +1; Fe3+, Fe = +3; O2-, O = -2 Oxidation numbers are hypothetical charges on atoms.

  7. Assigning Oxidation Numbers For ionic compounds, oxidation numbers can be assigned using the expected charges from the periodic table. The sum of the oxidation numbers of the atoms in the compound must equal 0.

  8. Assigning Oxidation Numbers For ionic compounds, oxidation numbers can be assigned using the expected charges from the periodic table. Alkali Metals = +1

  9. Assigning Oxidation Numbers For ionic compounds, oxidation numbers can be assigned using the expected charges from the periodic table. Alkaline Earth Metals = +2

  10. Assigning Oxidation Numbers For ionic compounds, oxidation numbers can be assigned using the expected charges from the periodic table. Group 13 Boron Group = +3

  11. Assigning Oxidation Numbers For ionic compounds, oxidation numbers can be assigned using the expected charges from the periodic table. Group 15 Nonmetals = -3

  12. Assigning Oxidation Numbers For ionic compounds, oxidation numbers can be assigned using the expected charges from the periodic table. Group 16 Nonmetals = -2

  13. Assigning Oxidation Numbers For ionic compounds, oxidation numbers can be assigned using the expected charges from the periodic table. Halogens = -1

  14. Assigning Oxidation Numbers For ionic compounds, oxidation numbers can be assigned using the expected charges from the periodic table. Transition Metals depend on anion

  15. Ionic Compounds Example: +2 -1 MgCl2

  16. Ionic Compounds Example: -2 +1 Na2O

  17. Ionic Compounds Example: We don’t know iron’s oxidation number from the periodic table since it is a transition metal. ?? -2 Fe2O3

  18. Ionic Compounds But since we know the compound is neutral, the oxidation numbers must add up to zero. Therefore, Fe has a +3 oxidation number in this compound. Example: +3 -2 Fe2O3

  19. Ionic Compounds Example: We don’t know chromium’s oxidation number from the periodic table since it is a transition metal. ?? -2 Cr2O72-

  20. Ionic Compounds But since this ion has a charge of -2, the oxidation numbers must add up to negative two. Therefore the oxidation number of Cr is +6 Example: +6 -2 Cr2O72-

  21. Covalent Compounds Covalent compounds are made of two nonmetals, which from the periodic table are always expected to be negative

  22. Covalent Compounds But since covalent compounds are neutral species, it is not possible for every element to retain its negative oxidation number

  23. Covalent Compounds ONLY THE MORE ELECTRONEGATIVE ELEMENT keeps its negative oxidation number. Other nonmetals must adapt to keep the compound neutral

  24. Electronegativity Trend Increases

  25. Covalent Compounds Example: -2 SO2 Since oxygen is the more electronegative element, it will have its normal oxidation number.

  26. Covalent Compounds Example: +4 -2 SO2 The compound is neutral, so the oxidation number of sulfur will be sufficient to balance out the two oxygen atoms.

  27. Covalent Compounds Example: -1 OF2 Since fluorine is the more electronegative element, it will have its normal oxidation number.

  28. Covalent Compounds Example: +2 -1 OF2 The compound is neutral, so the oxidation number of oxygen will be sufficient to balance out the two fluorine atoms.

  29. Covalent Compounds Example: -2 PO43- Since oxygen is the more electronegative element, it will have its normal oxidation number.

  30. Covalent Compounds Example: +5 -2 PO43- The ion has a charge of negative three, so the oxidation numbers must add up to the total charge of the ion. The sum of the oxidation numbers in the formula of a polyatomic ion is equal to its ionic charge

  31. Ionic Compounds with Polyatomic Example: +2 CaSO4 This is an ionic compound, so the charge of the metal cation is its oxidation number

  32. Ionic Compounds with Polyatomics Example: +2 CaSO4 The anion is a polyatomic ion, sulfate, and the charge of sulfate is negative two. So the oxidation numbers of sulfur and oxygen must add to -2

  33. Ionic Compounds with Polyatomics Example: +2 -2 CaSO4 Oxygen is the more electronegative of the two, so it keeps its normal oxidation number.

  34. Ionic Compounds with Polyatomics Example: +2 +6 -2 CaSO4 Sulfur and the four oxygen atoms must add to negative two (the charge of the sulfate anion).

  35. Ionic Compounds with Polyatomics Example: Pb(OH)4 This is an ionic compound, so the charge of the metal cation is its oxidation number. But this is a transition metal, so we cannot know it from its position on the periodic table.

  36. Ionic Compounds with Polyatomics Example: +4 Pb(OH)4 But the anion, the hydroxide ion, carries a charge of negative one. All four hydroxides are negative one, but since the compound is neutral, the oxidation number of lead must balance it out.

  37. Ionic Compounds with Polyatomics Example: +4 -2 Pb(OH)4 Within the anion, oxygen is the more electronegative of the two elements, and keeps its normal oxidation number.

  38. Ionic Compounds with Polyatomics Example: +4 -2 +1 Pb(OH)4 Within the hydroxide ion, the oxygen and hydrogen must add to the charge of the ion, -1

  39. Assign oxidation numbers to all of the elements: Li2O Li = +1 O = -2 PF3 P = +3 F = -1 HNO3 H = +1 +5 O = -2 N = MnO4- Mn = +7 O = -2 Cr2O72- Cr = +6 -2 O =

  40. What is the oxidation state of the highlighted element? P2O5 +5 diphosphorous pentoxide NaH -1 sodium hydride SnBr4 +4 tin (IV) bromide BaO2 -1 barium peroxide

  41. Identifying Redox Equations • In general, all chemical reactions can be assigned to one of two classes: • oxidation-reduction, in which electrons are transferred: • Single-replacement, combination, decomposition, and combustion • this second class has no electron transfer, and includes all others: • Double-replacement and acid-base reactions

  42. Identifying Redox Equations • In an electrical storm, nitrogen and oxygen react to form nitrogen monoxide: • N2(g) + O2(g)→ 2NO(g) • Is this a redox reaction? • If the oxidation number of an element in a reacting species changes, then that element has undergone either oxidation or reduction; therefore, the reaction as a whole must be a redox.

  43. A half-reaction is an equation showing just the oxidation or just the reduction that takes place • they are then balanced separately, and finally combined • Step 1: write unbalanced equation • Step 2: write separate half-reaction equations for oxidation and reduction • Step 3: balance all non H or O atoms in the half-reactions (e.g. Cr, Cu, Au, N)

  44. Step 4: add H2O to balance any O atoms • Step 5: add H+ to balance the H atoms • Step 6: add electrons to the most positive side of each half-reaction to balance the charges (remember one will be reduction, one will be oxidation) • Step 7: multiply each half-reaction by a number to make the electrons equal in both • Step 8: add the balanced half-reactions to show an overall equation

  45. Oxidizing and Reducing Agents • Now the confusing part… • CuO + H2 Cu + H2O • Cu goes from +2 to 0 • Cu is reduced, therefore it is called an oxidizing agent because it causes some other substance to be oxidized • H goes from 0 to +1 • H is oxidized, therefore it is called a reducing agent because it causes some other substance to be reduced.

  46. Some common oxidizing agents • Oxygen! Is an oxidant because it CAUSES oxidation, it itself is reduced though • Oxidized coal in electric power • Gas in automobiles • Wood in campfires • Food we eat • Antiseptics • Hydrogen Peroxide • Benzoyl peroxide • Disinfectants • Chlorine

  47. Some common reducing agents • Metals are REDUCTANTS! They cause REDUCTION, but themselves are oxidized (loss electrons) • Antioxidants • Ascorbic acid is used to prevent the browning of fruits by inhibiting air oxidation • Many antioxidants are believed to retard various oxidation reactions that are potentially damaging to vital components of living cells