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Oxidation Reduction Redox An introduction to electron transfer reactions

Oxidation Reduction Redox An introduction to electron transfer reactions. Oxidation-Reduction: A Reaction. Oxidation: When a substances loses electrons. Reduction: When a substance gains electrons. Consider: Ca (s) + 2H + (aq)  Ca 2+ (aq) + H 2(g) .

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Oxidation Reduction Redox An introduction to electron transfer reactions

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  1. Oxidation ReductionRedoxAn introduction to electron transfer reactions

  2. Oxidation-Reduction: A Reaction • Oxidation: When a substances loses electrons. • Reduction: When a substance gains electrons. • Consider: • Ca(s) + 2H+(aq)  Ca2+(aq) + H2(g). • The neutral Ca(s) has lost two e- to 2 H+ to become Ca2+. • We say Ca has been oxidized to Ca2+ • At the same time 2 electrons are gained by 2 H+ to form H2 . • We say H+ is reduced to H2 .

  3. Redox Reaction with Air • Consider the reaction of Ca with O2: • 2Ca(s) + O2(g)  2CaO(s) • Ca is easily oxidized in air. • On the left there is shiny Ca metal. • On the right we see a white powder – Calcium oxide. • Again, Ca(s) loses electrons and is oxidized to Ca+2 • And the neutral O2 has gained electrons from the Ca to become O2- in CaO. • We say O2 has been reduced to O2-.

  4. Electron Transfer and Terminology • Lose electrons: Oxidation • Gain electrons: Reduction. GER Leo

  5. It Takes Two: Oxidation-Reduction • In all reduction-oxidation (redox) reactions, one species is reduced at the same time as another is oxidized. • Oxidizing Agent: • The species which causes oxidation is called the oxidizing agent. The substance which is oxidized loses electrons to the other. • The oxidizing agent is always reduced • Reducing Agent: • The species which causes reduction is called the reducing agent. • The substance which is reduces gains electrons from the other. • The Reducing agent is always oxidized

  6. Oxidation of Metals with Acids • It is common for metal to produce hydrogen gas when they react with acids. For example, the reaction between Mg and HCl: • Mg(s) + 2HCI(aq)  MgCl2(aq) + H2(g) . • In this reaction, Mg is oxidized and H in HCl is reduced. • Note the change in oxidation state for these species: • Mg0 Mg+2 in MgCl2 • & • H+ in HCl  H0 in H2

  7. Redox reaction with Acid • It is possible for metals to be oxidized with salt: • Fe(s) + Ni(NO3)2 (aq)  Fe(NO3)2 (aq) + Ni (s) . • Molecular Equation • The net ionic equation shows the redox chemistry well: • Fe(s) + Ni+2(aq) Fe2+(aq) + Ni (s) • Net ionic Equation • In this reaction iron has been oxidized to Fe2+ while the Ni+2 has been reduced to Ni0. • What determines whether the reaction occurs ?

  8. The Activity Series • Metals can be placed in order of their tendencies for losing electrons. This is called the activity series.

  9. Competition For e- Transfer • Consider: Na, Mg, Al, • Metallic character decreases left to right. • Metal tend to give up electrons. • Now consider the reaction: • Na + AlCl3 ??? (NaCl + Al) • To determine if the reaction occurs, the question is to determine which metal has a greater affinity for electrons (or which is willing to lose e- ). • Na is more willing to lose e- than Al • Al is more willing to accept e- (less metallic) • Conclude: The reaction occurs. • 3Na + AlCl3 3NaCl + Al

  10. Reading Activity Table • A metal in the activity series can only be oxidized by a metal ion below it. • In our example, Na • is oxidized by Al. • The metals at the top of the activity series are called active metals. The metals at the bottom of the activity series are called noble metals.

  11. Example: Silver and Copper • If we place Cu into a solution of Ag+ ions, will copper plate out of solution ? • Cu(s) + 2AgNO3(aq)  ? [Cu(NO3)2(aq) + 2Ag (s)] • or Cu (s) + 2Ag+ (aq)  ? [Cu2+(aq) + 2Ag (s)] • Which metal is active? Which is noble ? • g CugAg • \Therefore, Cu 2+ ions is be formed because Cu is above Ag in the activity series. Copper Cu g Cu2+ + 2 e-Silver Ag g Ag+ + e-

  12. Example: Redox Reaction • B&L 4.47: Based on the activity series, what is the outcome of the following reaction ? • b) Ag(s) + PbNO3 (aq)  ? c) Cr (s) + NiSO4(aq)  ? • e) H2(g) + CuCl2(aq)  ? f) Ba (s) + H2O (l) ? • b) Ag vs. Pb , Pb is more active, no reaction occurs • c) Cr vs. Ni , Cr is more active, reaction occurs • Cr (s) + NiSO4(aq)  Ni (s) + CrSO4(aq) • d) H2 vs. Cu , H2 is more active, reaction occurs • H2(g) + CuCl2(aq)  2HCl(aq) + Cu (s) • e) Ba vs. H2 , Ba is more active, reaction occurs • Ba (s) + H2O (l) H2(g) + Ba(OH)2(aq)

  13. Summary • Redox - Oxidation/Reduction reaction • Oxidation-Lose electron (LEO) • Reduction-Gain electron (GER) • Activity Series-Table showing elements’ relative ease of oxidation. • Active MetalMetal which prefers to lose e- and therefore prefers the oxidized form. • Noble MetalMetal which do not lose e- and therefore prefers the zero state.

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