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REDOX REACTIONS

REDOX REACTIONS. Previously: What happened to oxygen when it reacted During reactions oxygen would take on electrons Now: When any element gains electrons. REDUCTION. OXIDATION. Previous: What happened to an element when it reacted with oxygen

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REDOX REACTIONS

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  1. REDOX REACTIONS

  2. Previously: What happened to oxygen when it reacted During reactions oxygen would take on electrons Now: When any element gains electrons REDUCTION

  3. OXIDATION • Previous: What happened to an element when it reacted with oxygen • During reaction, oxygen would take the electrons from elements • Now: When any element loses electrons

  4. REDOX • Reduction and oxidation NEVER happen by themselves • When an element is oxidized, there is another element that is reduced • When an element is reduced, there is an another element that is oxidized • LEO the Lion says GER • Lose an electronoxidized • Gain an electronsreduced

  5. OXIDATION NUMBERS • Oxidation numbers can be assigned using the periodic table if the compound is an ionic compound. • Covalent compounds are made of two nonmetals, which from the periodic table are always expected to be negative. • Since covalent compounds are neutral, it is not possible for every element to retain its negative oxidation number. • Only the more electronegative element stays negative; least electronegative element changes to positive • Oxidation number is different from formal charge. • Using oxidation number gives us another way to account for electrons in chemical changes.

  6. RULES FOR DETERMINING OXIDATION NUMBERS • The oxidation number of any uncombined element is zero • The oxidation number of a monatomic ion equals its charge • Oxygen’s oxidation number is -2, except in peroxides (H2O2) where it is -1 or when it bonds with fluorine where it will be +2 • The oxidation number of hydrogen is +1 except when it bonds with metals to from metal hydrides, where it is -1 or in the polyatomic ion NH4 where it is also -1 • The sum of the oxidation numbers for a compound must equal zero • The sum of the oxidation numbers in the formula of a polyatomic ion is equal to its charge.

  7. RULES FOR DETERMINING OXIDATION NUMBERS • 1-2: 2Na + Cl2 2NaCl • 3-4: H2O • 5: CaCl2 Ca(OH)2 • 6: NO3- SO4-2

  8. REDOX REACTIONS • Any reaction where the oxidation number of an element is different on the two sides of the chemical equation is a redox equation • Mg + 2HCl  MgCl2 + H2 • Oxidation = loss of electrons • Reduction = gaining electrons • You will represent redox reactions using half reactions

  9. REDOX REACTIONS • Half Reactions • Mg + 2HCl  MgCl2 + H2 • The total exchange of all electrons is accounted for in the combined half reactions

  10. AGENTS • Reducing Agent: An element that serves as the source of electrons is known as the reducing agent • The reducing agent is the one that is oxidized • Oxidizing Agent: An element that receives the electrons is known as the oxidizing agent • The oxidizing agent is the one that is reduced

  11. EXAMPLES • 4Fe + 3O2 2Fe2O3 • 2Fe2O3  4Fe + 3O2 • AgNO3 + Mg  Mg(NO3)2 + Ag

  12. REDOX REAL-LIFE EXAMPLES • Galvanic Wet Cell

  13. REDOX REAL-LIFE EXAMPLES • How does this work • If you bathe two different strips in a conductive solution, and connect them externally with a wire • Reactions between the electrodes and the solution furnish the circuit with charges continually. • The process that produces the electrical energy continues and becomes useful. • Spontaneously conversion of chemical energy to electrical energy. • The Copper (Cu) atoms attract electrons more than do the Zinc (Zn) atoms. • Zinc is more active and yields its electrons more easily.

  14. REDOX REAL-LIFE EXAMPLES

  15. REDOX REAL-LIFE EXAMPLES • AN ALKALINE BATTERY • Anode: Zn cap: • Zn(s) → Zn2+(aq) + 2e- • Cathode: MnO2, NH4Cl and carbon paste: 2 NH4 +(aq) + 2 MnO2(s) + 2e- → Mn2O3(s) + 2NH3(aq) + 2H2O(l) • Graphite rod in the center - inert cathode. • Alkaline battery, NH4Cl is replaced with KOH. • Anode: Zn powder mixed in a gel:

  16. REDOX REAL-LIFE EXAMPLES

  17. REDOX REAL-LIFE EXAMPLES • CORROSION OF IRON • Since E°(Fe2+/Fe) < E°(O2/H2O) iron can be oxidized by oxygen. • Cathode • O2(g) + 4H+(aq) + 4e- → 2H2O(l). • Anode • Fe(s) → Fe2+(aq) + 2e-. • Fe2+ initially formed – further oxidized to Fe3+ which forms rust, Fe2O3• xH2O(s).

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