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Electrochemistry Chapter 9 Kulkarni

Electrochemistry Chapter 9 Kulkarni. Electrochemistry: Study of chemical processes that involve the exchange of electrons. Reduction : Gain of electrons. The substance that gains electrons is said to be reduced Cl 2 + 2e - 2Cl - The process is reduction

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Electrochemistry Chapter 9 Kulkarni

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  1. ElectrochemistryChapter 9Kulkarni

  2. Electrochemistry: Study of chemical processes that involve the exchange of electrons.

  3. Reduction: Gain of electrons. The substance that gains electrons is said to be reduced Cl2+ 2e- 2Cl- The process is reduction Chlorine is reduced to chloride.

  4. Oxidation: Loss of electrons. The substance that loses electrons is said to be oxidized Ca 2e- + Ca+2 The process is oxidation Calcium is oxidized to calcium ion.

  5. Redox Oxidation cannot occur without reduction and vice versa. That’s why they are called “redox” reactions! Ca + Cl2 CaCl2

  6. “LEO the lion says GER” Loss Gain Electrons Electrons Oxidation Reduction

  7. “OIL RIG” Oxidation Is Loss Reduction Is Gain

  8. Oxidation Number • Used to keep track of electrons • Shows the general distribution of electrons NOT absolute charge!

  9. RULES! • Uncombined elements: Oxid. # = zero. Ex. Cu, Cl2, P

  10. Binary ionic compounds: Oxid. # = ionic charge Ex. CaCl2: Ca+2 o.#. Is +2 and Cl-1 o.#. is -1

  11. 3. Metals in compounds: Group 1 +1 Group 2 +2 Aluminum +3 Zn and Cd +2 Silver +1 Transitions ionic charge

  12. Fluorine: -1 always • Hydrogen: • Usually +1 • With metals -1 Ex. NaH, CaH2 (hydrides)

  13. Oxygen Usually -2 In peroxides -1 Ex: H2O2, group I oxides (Li2O2) With Fluorine +2 (OF2)

  14. 7. The sum of oxidation numbers in a neutral compound is zero. CaCl2: (+2) + 2(-1) = 0

  15. Covalent Compounds The more electronegative element is assigned all of the electrons as if it were ionic. Ex. NO2 – Oxygens are each -2 so the nitrogen is +4. (x) + 2(-2) = 0, x = +4

  16. The sum of oxidation numbers for a polyatomic ion equals the charge on the polyatomic ion. PO4-3: (+5) + 4(-2) = -3

  17. PRACTICE!

  18. MnO2 + 4HCl MnCl2 + 2H2O +Cl2 Redox Reactions In redox reactions, one or more atoms change oxidation numbers. +4 -2 +1 -1 +2 -1 +1 -2 0 Mn+4 is reduced Chloride (Cl-) is oxidized

  19. Half-Reactions Two parts of a redox reaction written separately Mg + Cl2 MgCl2 • Mg Mg+2 + 2e- • Cl2 + 2e- 2Cl- Oxid. Red.

  20. Reducing Agent: Helps another element become reduced. (it is oxidized in the process) Oxidizing Agent: Helps another element become oxidized. (it is reduced in the process)

  21. NOTE: When discussing what is reduced/oxidized or what is the oxidizing/reducing agents we are always referring to REACTANTS with their OXIDATION NUMBER!

  22. PRACTICE!

  23. Balancing Redox – ½ rxn method • Write down the chemical equation. • Determine all oxidation numbers. • Write down balanced ½ rxns.

  24. Find a common multiple for the number of electrons in each ½ rxn such that: # e-’s gained = # e-’s lost • Move coefficients from the ½ rxn to the complete rxn.

  25. Complete the balancing process for the remaining elements. NOTE: If not all of an element is oxidized or reduced you may not be able to use the coefficient.

  26. PRACTICE!

  27. Balancing in Acidic Soln: • Assign Oxidation #s • Write the ½ reactions as before

  28. Balance the charge with the addition of e-’s Find a common multiple for the number of electrons in each ½ rxn such that: # e-’s gained = # e-’s lost

  29. Place coefficients into the original reaction. • Balance the oxygens with water. • Balance the hydrogens with H+. • Cancel of necessary

  30. PRACTICE!

  31. Balancing in Basic Soln: • Balance as if it were Acidic Solution. • Find a common multiple for the number of electrons in each ½ rxn such that: # e-’s gained = # e-’s lost

  32. Neutralize the H+ with OH-. Add the same amount of OH- to the opposite side of the equation.

  33. Cancel as needed H2O’s Hint: H+ + OH- = H2O That is why we neutralize the H+ in step number 2.

  34. PRACTICE!

  35. PRACTICE! In Class WKS

  36. Electrochemical Cells System of two cells in which chemical energy is converted to electrical energy or vice versa.

  37. 1. Galvanic/Voltaic Cell SPONTANEOUS electro-chemical cell. The transfer of electrons produces energy Ex. batteries

  38. 2. Electrolytic Cells NON-SPONTANEOUS – energy must be added to bring about a chemical change. Ex. Electroplating

  39. 3. Components • ELECTRODES – conductor that establishes electrical contact with a nonmetallic, electrolytic part of the cell.

  40. ANODE – where oxidation takes place. CATHODE – where reduction takes place “An ox and a red cat”

  41. Figure 18.1: Schematic for separating the oxidizing and reducing agents.

  42. Figure 18.2: Electron flow.

  43. Figure 18.3: Ion flow keeps the charge neutral.

  44. POROUS BARRIER/ SALT BRIDGE Prevents half-reactions from mixing but allows ions to move freely to balance charges in both cells.

  45. Figure 18.4: The salt bridge contains a strong electrolyte.

  46. Figure 18.4: The porous disk allows ion flow.

  47. Figure 18.5: Schematic of a galvanic cell. Click on picture to go to http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/galvan5.swf

  48. Electrolytic Cell – requires a power supply

  49. Electroplating – Electrolytic cell

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