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Chapter 14 – Fundamentals of Electrochemistry. Homework - Due Friday, April 1 Problems: 14-4, 14-5, 14-8, 14-12, 14-15, 14-17, 14-18, 14-25, 14-26, 14-41, . Electrochemistry Review of the Basics. Oxidation Loss of electrons Always occurs at the anode

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Chapter 14 – Fundamentals of Electrochemistry

Homework - Due Friday, April 1

Problems: 14-4, 14-5, 14-8, 14-12, 14-15, 14-17, 14-18, 14-25, 14-26, 14-41,

CHM 320 - Lecture 23 Chapt 14

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Electrochemistry

Review of the Basics

  • Oxidation
    • Loss of electrons
    • Always occurs at the anode
    • Happens because of the action of a reducing agent
  • Reduction
    • Gain of electrons (charge is reduced)
    • Always occurs at the cathode
    • Happens because of the action of the oxidizing agent

Fe+3 + e- = Fe+2

CHM 320 - Lecture 23 Chapt 14

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Redox Reaction

ox1 + red2 = red1 + ox2

Example:

2(e- + Fe+3 = Fe+2)

+ Sn+2 = Sn+4 + 2 e-

-------------------------------------------

2 Fe+3 + Sn+2 = 2 Fe+2 + Sn+4

So… ox1= Fe+3 , red2 = Sn+2, red1=Fe+2, ox2=Sn+4

CHM 320 - Lecture 23 Chapt 14

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4

Electric Charge (in Coulombs) and Work

  • Voltage represents electrical potential (potential to do work)
  • If some total charge in coulombs (q) is moved through some electrical potential (E, in volts V) then work is done!
  • The charge in coulombs (q) is equal to the number of moles of electrons (n) times the Faraday Constant (F)

CHM 320 - Lecture 23 Chapt 14

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Ohm’s Law and Power

  • Ohm’s law relates electrical resistance, current and potential!
  • Power is the work done in some unit time (e.g. joules of work per second)
  • The units of Power are Watts (W)
  • Ohm’s law and power are related!

CHM 320 - Lecture 23 Chapt 14

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Let’s Work Some Problems

  • A 6.00V battery is connected across a 2.00 K resistor, how many electrons flow through the circuit per second?
  • How many joules of heat (heat is work) are produced per electron?
  • What voltage would the battery need to be to deliver a power at 100.0 Watts?

CHM 320 - Lecture 23 Chapt 14

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Electrochemical Cells

  • There are two basic electrochemical cells:
    • A GALVANIC cell uses spontaneous chemical reactions to generate electricity
    • A ELECTROYLTIC cell requires an electrical potential to be applied to the cell to drive some reaction.
  • A complete cell contains:
    • anode
    • cathode
    • completed circuit (for electrons to flow)
    • a salt bridge (usually!)
    • an electrolyte solution
    • chemical species that undergo reaction.

CHM 320 - Lecture 23 Chapt 14

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Galvanic Cell

Cells and Cell Reactions

Overall Cell Reaction

Zn(s) + Cu+2(aq) ---> Zn+2(aq) + Cu(s)

oxidation half reaction

anode Zn(s) ---> Zn+2(aq) + 2 e-

reduction half reaction

cathode Cu+2(aq) + 2 e- ---> Cu(s)

CHM 320 - Lecture 23 Chapt 14

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The Standard Hydrogen Electrode (SHE)

  • The basis by which all other measurements are made.
  • Assigned a potential of zero by definition!
  • Not practical for regular use

Hydrogen Half-Cell

H2(g) = 2 H+(aq) + 2 e-

reversible reaction

SHE consists of a platinum electrode covered with a fine powder of platinum around which H2(g) is bubbled. Its potential is defined as zero volts.

CHM 320 - Lecture 23 Chapt 14

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Standard Potentials

  • Standardized potentials (Eo), listed as reductions, for all half-reactions
  • Measured versus the S.H.E (0)
  • Used in predicting the action in either a galvanic cell or how much energy would be needed to force a specific reaction in a non-spontaneous cell
  • Assumes an activity of one for the species of interest (usually a fair approximation) at a known temperature in a cell with the S.H.E.
  • Assumes that the cell of interest is connected to the (+) terminal of the potentiometer (voltmeter) and the S.H.E. is connected to the (-) terminal

CHM 320 - Lecture 23 Chapt 14

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Better Oxidizing Agents

in upper left hand corner.

Better Reducing

Agents in lower

Right hand corner

CHM 320 - Lecture 23 Chapt 14