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REDOX REACTIONS

Electrochemical Reactions. REDOX REACTIONS. REACTIONS. Day 1 Review Oxidation numbers. Electron Transfer Reactions. Electron transfer reactions are oxidation-reduction or redox reactions.

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REDOX REACTIONS

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  1. Electrochemical Reactions REDOX REACTIONS REACTIONS Day 1 Review Oxidation numbers

  2. Electron Transfer Reactions • Electron transfer reactions are oxidation-reduction or redox reactions. • Redox reactions can result in the generation of an electric current or be caused by imposing an electric current. • Therefore, this field of chemistry is often called ELECTROCHEMISTRY.

  3. Why Study Redox Reactions Batteries Corrosion Manufacturing metals Fuels

  4. Net Ionic Equations Mg(s) + 2 HCl(aq) --> H2(g) + MgCl2(aq) We really should write Mg(s) + 2 H+(aq) + 2 Cl-(aq) ---> H2(g) + Mg2+(aq) + 2 Cl-(aq) The two Cl- ions are SPECTATOR IONS — they do not participate. Could have used NO3-.

  5. Net Ionic Equations Mg(s) + 2 HCl(aq) --> H2(g) + MgCl2(aq) Mg(s) + 2 H+(aq) + 2 Cl-(aq) ---> H2(g) + Mg2+(aq) + 2 Cl-(aq) We leave the spectator ions out — Mg(s) + 2 H+(aq) ---> H2(g) + Mg2+(aq) to give the NET IONIC EQUATION

  6. Water Solubility of Ionic Compounds If one ion from the “Soluble Compound” list is present, the compound is water soluble.

  7. Insoluble products: Precipitation formation

  8. K+(aq) + MnO4-(aq) IONIC COMPOUNDSCompounds in Aqueous Solution Most redox reactions are in water — aqueous solutions - and involve ionic compounds KMnO4 in water

  9. Aqueous Solutions Solutions that conduct electricity are called ELECTROLYTES Strong electrolytesdissociate completely (or nearly so) into ions.

  10. Aqueous Solutions Acids that ionize only to a small extent are called weak electrolyte. CH3CO2H(aq) ---> CH3CO2-(aq) + H+(aq)

  11. Aqueous Solutions Some compounds dissolve in water but do not conduct electricity. They are called nonelectrolytes. Examples include: sugar ethanol ethylene glycol

  12. Know the strong acids & bases!

  13. HNO3 ACIDS An acid -------> H+ in water Some strongacids are HCl hydrochloric H2SO4 sulfuric HClO4perchloric HNO3 nitric

  14. BASESsee Screen 5.9 and Table 5.2 Base ---> OH- in water NaOH(aq) ---> Na+(aq) + OH-(aq) NaOH is a strong base

  15. Oxidation/Reduction Reactions Thermite reaction Fe2O3(s) + 2 Al(s) ----> 2 Fe(s) + Al2O3(s)

  16. REDOX REACTIONS Oxidation— lose of electrons Reduction— gaining of electrons MUST always occur together 2 H2(g) + O2(g) g 2 H2O(liq) Mg(s) + 2 HCl(aq) g MgCl2(aq) + H2(g) 2 Al(s) + 3 Cu2+(aq) g 2 Al3+(aq) + 3 Cu(s) Fe2O3(s) + 2 Al(s) g 2 Fe(s) + Al2O3(s)

  17. REDOX REACTIONS Cu(s) + 2 Ag+(aq) g Cu2+(aq) + 2 Ag(s) In all reactions if something has been oxidized then something has also been reduced

  18. REDOX REACTIONS Redox reactions are characterized by ELECTRON TRANSFER between an electron donor and electron acceptor. Transfer leads to— 1. increase in oxidation number of some element = OXIDATION 2. decrease in oxidation number of some element = REDUCTION

  19. OXIDATION NUMBERS The electric charge an element APPEARS to have when electrons are counted by some arbitrary rules: 1. An element has ox. no. = 0. Zn O2 I2 S8 2. In simple ions, ox. no. = charge on ion. -1 for Cl- +2 for Mg2+

  20. OXIDATION NUMBERS 3. In a compound: O has ox. no. = -2 (except in peroxides: in H2O2, O = -1) 4. In a Molecule: Ox. no. of H = +1 (except when H is associated with a metal as in NaH where it is -1) 5. Algebraic sum of oxidation numbers = 0 for a compound = overall charge for an ion

  21. Oxidation number of F in HF? OXIDATION NUMBERS NH3 N = ClO- Cl = H3PO4 P = MnO4- Mn = Cr2O72- Cr = C3H8 C =

  22. Recognizing a Redox Reaction Corrosion of aluminum 2 Al(s) + 3 Cu2+(aq) --> 2 Al3+(aq) + 3 Cu(s) Al(s) --> Al3+(aq) + 3 e- • Ox. no. of Al increases as e- are donated by the metal. • Therefore, Al is OXIDIZED • Al is the REDUCING AGENT in this balanced half-reaction.

  23. Recognizing a Redox Reaction Corrosion of aluminum 2 Al(s) + 3 Cu2+(aq) --> 2 Al3+(aq) + 3 Cu(s) Cu2+(aq) + 2 e- --> Cu(s) • Ox. no. of Cu decreases as e- are accepted by the ion. • Therefore, Cu is REDUCED • Cu is the OXIDIZING AGENT in this balanced half-reaction.

  24. Common Oxidizing and Reducing Agents

  25. Metals (Cu) are “loser”: oxidized Metals (K) are reducing agents HNO3 is the oxidizing agent Common Oxidizing and Reducing Agents hydrogen is reduced Cu + HNO3 --> Cu2+ + NO2 2 K + 2 H2O --> 2 KOH + H2

  26. Examples of Redox Reactions Nonmetal (S) + Oxygen Metal (Mg) + Oxygen

  27. Examples of Redox Reactions Metal + acid Zn + HCl Zn = reducing agent H+ = oxidizing agent

  28. Examples of Redox Reactions Metal + halogen 2 Al + 3 Br2 ---> Al2Br6

  29. Metal + acid Cu + HNO3 Cu = reducing agent HNO3 = oxidizing agent

  30. Balancing Equations Cu + Ag+ --give--> Cu2+ + Ag Need to Balance BOTH mass and CHARGE Step 1:Divide into half-reactions: one for oxidation and the other for reduction. Ox Cu ---> Cu2+ Red Ag+ ---> Ag

  31. Balancing Equations Step 2: Balance each for mass. Already done in this case. Step 3: Balance each half-reaction for charge by adding electrons. Ox Cu ---> Cu2+ + 2e- Red Ag+ + e- ---> Ag

  32. Balancing Equations Step 4:Multiply each half-reaction by a factor to have the electrons lost equal to number gained Cu ---> Cu2+ + 2e- 2 Ag+ + 2 e- ---> 2 Ag Step 5: Add to give the overall equation. Cu + 2 Ag+ ---> Cu2+ + 2Ag The equation is now balanced for BOTH charge and mass.

  33. Review of Terminology for Redox Reactions • OXIDATION—loss of electron(s) by a species; increase in oxidation number. • REDUCTION—gain of electron(s); decrease in oxidation number. • OXIDIZING AGENT—electron acceptor; species is reduced. • REDUCING AGENT—electron donor; species is oxidized.

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