Introduction to Oxidation-Reduction Reactions

1. Introduction to Oxidation-Reduction Reactions Electron Transfer Reactions

2. Types of Chemical Reactions • There are four types of chemical reactions: • Acid/Base • Precipitation/Solubility • Complex Formation/Complex Dissociation • Oxidation/Reduction • Any chemical reaction consists of one (or more) of these basic categories.

3. Oxidation/Reduction Reactions • Acid/Base reactions • involve a donation /acceptance of protons • Precipitation/ Solubility reactions • involve a donation/ acceptance of negative charge what is being donated and accepted in a redox reaction?

4. Oxidation/Reduction Reactions • Electrons! • Consider the reaction taking place in a disposable battery: 2Zn + 3MnO2 Mn3O4 + 2ZnOHow can you tell that electrons are being donated and accepted? Which species is donating electron( s) and which is accepting electron (s)?

5. REDOX REACTIONS Redox reactions are characterized by ELECTRON TRANSFER between an electron donor and electron acceptor.

6. REDOX REACTIONS Transfer leads to— • increase in oxidation number of some element = OXIDATION 2. decrease in oxidation number of some element = REDUCTION

7. Electron Transfer in Redox Reactions • Oxidation • Loss of electrons • Gain in oxygen • Reduction • Gain of electrons • Loss of oxygen • “LEO the lion goes Ger”

8. Example • The reaction of a metal and non-metal • All the electrons must be accounted for! Mg S Mg 2+ S2- + → +

9. Oxidation-Reduction • Oxidation means an increase in oxidation state - lose electrons. • Reduction means a decrease in oxidation state - gain electrons. • The substance that is oxidized is called the reducing agent. • The substance that is reduced is called the oxidizing agent.

10. Assigning Oxidation States • An Oxidation-reduction reaction involves the transfer of electrons. • You should memorize these rules

11. Rules for Oxidation States • The charge the atom would have in a molecule (or an ionic compound) if electrons were completely transferred. • The oxidation state of elements in their standard states is zero. • Example: Na, Be, K, Pb, H2, O2, P4 = 0

12. Assigning Oxidation States • Oxidation state for monatomic ions are the same as their charge. • Example: Li+, Li = +1; Fe3+, Fe = +3; O2-, O = -2 • Oxygen is assigned an oxidation state of -2 in its covalent compounds except as a peroxide.

13. Rules for Oxidation States • The oxidation number of hydrogen is +1except when it is bonded to metals in binary compounds. In these cases, its oxidation number is –1. • Group IA metals are +1, IIA metals are +2 and fluorine is always –1. • 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.

14. Practice in Oxidation States • Assign the oxidation states to each element in the following. • K2SO4 • NO3- • H2SO4 • Fe2O3 • Fe3O4

15. Identify the • Oxidizing agent • Reducing agent • Substance oxidized • Substance reduced • On the worksheet

16. A + B C S + O2 SO2 C A + B 2KClO3 2KCl + 3O2 Types of Oxidation-Reduction Reactions Combination Reaction +4 -2 0 0 Decomposition Reaction +1 +5 -2 +1 -1 0

17. A + BC AC + B Sr + 2H2O Sr(OH)2 + H2 TiCl4 + 2Mg Ti + 2MgCl2 Cl2 + 2KBr 2KCl + Br2 Types of Oxidation-Reduction Reactions Displacement Reaction a.k.a Single Replacement +1 +2 0 0 Hydrogen Displacement +4 0 0 +2 Metal Displacement -1 0 0 -1 Halogen Displacement

18. M + BC AC + B Ca + 2H2O Ca(OH)2 + H2 The Activity Series for Metals Hydrogen Displacement Reaction M is metal BC is acid or H2O B is H2

19. Copper Demonstration • Copper Pennies reacting with nitric acid. • Can you figure out the equation?

20. Cl2 + 2OH- ClO- + Cl- + H2O Chlorine Chemistry Types of Oxidation-Reduction Reactions Disproportionation Reaction Element is simultaneously oxidized and reduced. +1 -1 0

21. Classify the following reactions. Ca2+ + CO32- CaCO3 NH3 + H+ NH4+ Zn + 2HCl ZnCl2 + H2 Ca + F2 CaF2

22. Half-Reactions • All redox reactions can be thought of as happening in two halves. • One produces electrons - Oxidation half. • The other requires electrons - Reduction half.

23. Half-Reactions • Write the half reactions for the following. • Na + Cl2→ Na+ + Cl- • SO3- + H+ + MnO4- → SO4- + H2O + Mn+2

24. Balancing Redox Equations • In aqueous solutions the key is the number of electrons produced must be the same as those required. • For reactions in acidic solution an 8 step procedure.

25. Balancing Redox Equations • Write separate half reactions • For each half reaction balance all reactants except H and O • Balance O using H2O

26. Acidic Solution • Balance H using H+ • Balance charge using e-

27. Acidic Solution • Multiply equations to make electrons equal • Add equations and cancel identical species • Check that charges and elements are balanced.

28. Practice • Balance the following reactions: • Sn 2+ (aq) + 2Fe 3+ → Sn 4+ (aq) + 2Fe 2+ • MnO4- (aq) + C2O4-2 (aq) → Mn2+ (aq) + CO2 (g)

29. Practice • The following reactions occur in aqueous solution. Balance them • Cr(OH)3 + OCl- + OH-® CrO4-2 + Cl- + H2O • MnO4- + Fe+2® Mn+2 + Fe+3

30. Now for a tough one • Fe(CN)6-4 + MnO4-® Mn+2 + Fe+3 +CO2 + NO3-

31. Basic Solution • Do everything you would with acid, but add one more step. • Add enough OH- to both sides to neutralize the H+ • CrI3 + Cl2® CrO4-+ IO4- + Cl- • CN- + MnO4- → CNO- + MnO2

32. Redox Titrations • Same as any other titration. • the permanganate ion is used often because it is its own indicator. MnO4- is purple, Mn+2 is colorless. When reaction solution remains clear, MnO4- is gone. • Chromate ion is also useful, but color change, orangish yellow to green, is harder to detect.