1. Oxidation – Reduction a.k.a. REDOX Textbook Sections: 4.4-4.6 and 20.1-20.2

2. Good website: http://www.wfu.edu/~ylwong/redox/

3. Oxidation – the loss of one or more electrons by a substance (element, compound, ion) Reduction – the gain of one or more electrons by a substance (element, compound, ion) O. I. L. R. I. G. O xidation I s L oss, R eduction I s G ain (of electrons) L.E.O. the lion says G.E.R. L oss of E lectrons is O xidation G ain of E lectrons is R eduction

5. Oxidation ---------> A -2  A -1 + electron A -1  A + electron A  A +1 + electron A +1  A +2 + electron <--------- Reduction

6. Redox reaction – a process where electrons are transferred from one substance to another How can you tell when a redox reaction is taking place? 1)Assign oxidation numbers to atoms in substances 2)Compare oxidation numbers before and after reaction to determine if atom has lost or gained electrons

7. Rules For Assigning Oxidation Numbers Use the Rules in Order 1.The oxidation number of an atom in an element is 0. Examples: Na, H 2, Br 2, S 8, Ne Ox. # 0 0 0 0 0 2.The oxidation number of a monatomic ion is the same as its charge. Examples: Na +1, Ca +2, Al +3, Cl -1, O -2 Ox # +1 +2 +3 -1 -2

8. 3.The sum of the oxidation numbers of all atoms in a neutral compound is zero. 4.The sum of the oxidation numbers of all atoms in an ion is equal to the charge on the ion.

9. 5. In compounds, fluorine is always assigned an oxidation number of -1 (the most electronegative element in a compound always has a negative oxidation number.) 6. Hydrogen’s oxidation number will be - +1 when bonded to a nonmetal (HCl) --1 when bonded to a metal (NaH) Examples: NaH CaH 2 HClH 2 S Na—H H—Ca—H H—Cl H—S—H +1 +2 +1 +1-2 +1

10. 7.Oxygen usually has an oxidation number of -2. Exceptions: in peroxides, oxygen will be -1 and when combined with only F, it will be +2 and in O 2 it will be 0 Examples: H 2 O CaO H 2 O 2 H — O — H Ca—O H—O—O—H O 2 -2 OF 2 [O—O] -2 F—O—F +1-2 +1 +2 -2 +1 +1 +2

11. 8.Halogens usually have an oxidation number of -1. Examples: NaCl MgI 2 OCl 2 HOBr Na—Cl, I—Mg—I, Cl—O—Cl, H—O—Br +1 +2 +1 -2 +1 -2 +1 ** If none of the above rules help you get started…look for a atom with a known charge and use that charge as its oxidation number CdS: Cd-S +2 -2

12. Use algebra to determine oxidation numbers of "difficult" atoms. Example:H 2 SO 4 H is +1 * 2 = +2 O is -2 * 4 = -8 2 + S + -8 = 0 S is +6 Example:ClO 4 -1 O is -2 * 4 = -8 -8 + Cl = -1 Cl is +7 Example:NH 4 +1 H is +1* 4 = +4 4 + N = 1 N is -3

13. FeSO 4 O is -2 *4 = -8 Recognize this as an ionic compound – sulfate has a -2 charge For sulfate: x + -8 = -2 S = +6 Then look at the compound as a whole Fe + 6 + -8 = 0 Fe = +2 C3H8C3H8 H is +1 * 8 = 8 3C + 8 = 0 3C = -8 C = - 8/3 oxidation numbers do NOT have to be integers

14. Once oxidation numbers have been assigned, compare them before and after the reaction. 4 Fe (s) + 3 O 2 (g)  2 Fe 2 O 3 (s) 0 0 +3 -2 Fe is oxidized, going from 0 to +3 O is reduced, going from 0 to -2 Notice that a total of 12 electrons were lost and 12 electrons were gained

15. 2 Fe 2 O 3 (s) + 3 C(s)  4 Fe(s) + 3 CO 2 (g) +3 -2 0 0 +4 -2 Fe is reduced, going from +3 to 0 C is oxidized, going from 0 to +4 O undergoes no change 12 electrons lost and 12 electrons gained As seen in the above examples, oxidation and reduction ALWAYS occur together

16. Reducing agent (reductant) causes reduction loses electrons undergoes oxidation oxidation number increases Oxidizing agent (oxidant) causes oxidation gains electrons undergoes reduction oxidation number decreases

17. Assign oxidation numbers, indicate what is oxidized and reduced, indicate what is the oxidizing agent and reducing agent Ca(s) + 2 H +1 (aq)  Ca +2 (aq) + H 2 (g) 0 +1 +2 0 Ca is oxidized – increasing from 0 to +2 H +1 is reduced – decreasing from +1 to 0 Ca is the reducing agent H +1 is the oxidizing agent

18. 2 Fe +2 (aq) + Cl 2 (aq)  2 Fe +3 (aq) + 2 Cl -1 (aq) +2 0 +3 -1 Fe +2 is oxidized – increasing from +2 to +3 Cl 2 is reduced – decreasing from 0 to -1 Fe +2 is the reducing agent Cl 2 is the oxidizing agent

19. In general, metals and anions act as reducing agents (are oxidized) and nonmetals and cations act a oxidizing agents (are reduced). Periodic table – in general, metals on left of table are more active, metals become less active as you move to the right side of the table

20. Predicting Products of Redox Reactions The simple ones you already know: Decomposition (except of acids, bases, carbonates & hydrates) Composition (except of two oxides) Combustion Replacement

21. Replacement Reactions Replacement Reactions are redox reactions. General pattern: A + B X  A X + B Mg(s) + 2 HCl(aq)  MgCl 2 (aq) + H 2 (g) The Mg is oxidized and the H + is reduced.

22. Fe(s) + Ni(NO 3 ) 2 (aq)  Fe(NO 3 ) 2 (aq) + Ni(s) The net ionic equation shows the redox chemistry well: Fe(s) + Ni +2 (aq)  Fe +2 (aq) + Ni(s) Fe is oxidized to Fe +2 Ni +2 is reduced to Ni. Always keep in mind that whenever one substance is oxidized, some other substance must be reduced.

23. The Activity Series The activity series is a list of metals in order of decreasing ease of oxidation. The metals at the top of the activity series are called active metals and are easily oxidized (they WANT to be ions) The metals at the bottom of the activity series are called noble metals and NOT easily oxidized (they WANT to be atoms)

24. Oxidation of copper metal by silver ions. (Dime Lab)

25. A metal in the activity series can ONLY be oxidized by a metal ion below it (metal doing the replacing has to be above what it is replacing in the activity series) If we place Cu into a solution of Ag + ions, then Cu +2 ions can be formed because Cu is above Ag in the activity series: Cu(s)+2AgNO 3 (aq)  Cu(NO 3 ) 2 (aq)+2Ag(s) or Cu(s) + 2Ag + (aq)  Cu +2 (aq) + 2Ag(s) This is only part of the story – more later !

26. Special Cases: 1. Hydrogen reacts with a hot metallic oxide to produce the metal element and water. Ex: H 2 + MgO  Mg + H 2 O 2. A metal sulfide reacts with oxygen to produce the metallic oxide and sulfur dioxide. Ex: 2MgS + 3O 2  2MgO + 2SO 2

27. 3. Chlorine gas reacts with dilute sodium hydroxide to produce sodium hypochlorite, sodium chloride, and water. Cl 2 + 2NaOH  NaClO + NaCl + H 2 O 4. Copper reacts with concentrated sulfuric acid to produce copper(II) sulfate, sulfur dioxide, and water. Cu + 2H 2 SO 4  CuSO 4 + SO 2 + 2H 2 O

28. 5. Copper reacts with dilute nitric acid to produce copper(II) nitrate, nitrogen monoxide, and water. Cu + HNO 3  Cu(NO 3 ) 2 + NO + H 2 O 6. Copper reacts with concentrated nitric acid to produce copper(II) nitrate, nitrogen dioxide, and water. Cu + HNO 3  Cu(NO 3 ) 2 + NO 2 + H 2 O

29. Reactants (oxidizing agents)Products MnO 4 - in acidic solutionMn 2+ MnO 2 in acidic solutionMn 2+ MnO 4 - in neutral or basic solutionMnO 2 (s) MnO 4 - in very basic solutionMnO 4 2- Cr 2 O 7 2- in acidic solutionCr 3+ HNO 3, concentratedNO 2 HNO 3, diluteNO H 2 SO 4, hot, concentratedSO 2 Metallic ion (higher charge)Metallous ion (lower charge) Elemental HalogenHalogen ion Na 2 O 2 NaOH HClO 4 Cl - C 2 O 4 2- CO 2 H2O2H2O2 H2OH2O Memorize… These are reduction reactions (oxidation number decreases)

30. Reactants (Reducing Agents)Products Halogen ionsHalogen element Metal elementMetal ion SO 3 2- or SO 2 SO 4 2- NO 2 - NO 3 - Halogen element, dilute basic solutionHypo-halogen-ite ion (Ex: ClO -, BrO - ) Halogen element, concentrated basic solution Halogen-ate ion (Ex: ClO 3 -, BrO 3 - ) Metallous ion (lower charge)Metallic ion (higher charge) H2O2H2O2 O2O2 Memorize… These are oxidation reactions (oxidation number increases)