Exp' 32 Galvanic Cells, the Nernst Equation p' 357

1
Exp. 32 Galvanic Cells, the Nernst Equation
p. 357

• To measure the relative reduction potentials for
  a number of redox couples
• To develop an understanding of the movement of
  electrons, anions, and cations in a galvanic
  cell.
• To study factors affecting cell potentials.
• To estimate the concentration of ions in solution
  using the Nernst equation.

2
Exp. 32 Introduction

• When iron corrodes, a change in the oxidation
  number of the iron atoms occurs.
• When gasoline burns, a change in the oxidation
  number of the carbon atoms occurs.
• A change in the oxidation number of an atom in a
  chemical reaction is the result of an exchange of
  electrons between the reactants.

3
Exp. 32 Introduction

• A chemical reaction that involves the
• transfer of electrons from one substance to
• another is an oxidation-reduction (redox)
• reaction.
• If you place copper wire into a silver ion
• solution, copper atoms spontaneously loose
• electrons (are oxidized) to the silver ions
• which are reduced.

4
Exp. 32 Introduction

• Silver ions migrate to the copper atoms to
• pick up electrons and form silver atoms at
• the copper metal/solution interface (picture,
• p. 357). The reaction that occurs at the
• interface is
• Cu(s) 2Ag(aq) ? 2Ag(s) Cu2(aq)

5
Exp. 32 Introduction

• Redox reactions can be divided into
• oxidation and reduction half-reactions.
• Each half-reaction, called a redox couple,
• consists of the reduced state and the
• oxidized state of the substance
• Cu(s) 2Ag(aq) ? 2Ag(s) Cu2(aq)
• Cu(s) ? Cu2(aq) 2e- oxidation ½ rxn
• 2 Ag(aq) 2e- ? 2 Ag(s) reduction ½ rxn

6
Exp. 32 Introduction

• A galvanic cell uses this spontaneous
• transfer of electrons. Instead of electrons
• being transferred at the interface between
• copper metal and silver ions, a galvanic
• cell separates the copper metal from the
• silver ions, forcing the electrons to pass
• externally through a wire.

7
Exp. 32 Introduction

• The two redox couples are placed in separate
• compartments called half-cells (fig. 32.1).
• Each half-cell consists of an electrode, usu.
• the metal of the redox couple and a solution
• containing the corresponding cation of the
• redox couple.
• The electrodes are connected by a wire and
• the solutions are connected by a salt bridge.

8
Exp. 32 Introduction

• The electrode at which reduction occurs is
• called the cathode ().
• The electrode at which oxidation occurs is
• called the anode (-)
• Electrons flow from the anode to the cathode.

9
Exp. 32 Introduction

• Different metals have different tendencies
• to oxidize similarly, their ions have
• different tendencies to undergo reduction.
• Cell potential (Ecell) the difference in
• tendencies of the two metals to oxidize
• (lose electrons) or of their ions to reduce
• (gain electrons).

10
Exp. 32 Introduction

• Reduction potential a common
• measurement of the tendency for a
• substance to gain electrons or, the value
• used to identify the relative ease of
• reduction for a half-reaction.
• Potentiometer a gauge that measures
  
• the Ecell

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Exp. 32 Introduction

• Ecell EAg,Ag ECu2,Cu
• The higher the reduction potential (more
• positive), the greater the tendency to
• undergo reduction.
• In our current example, this is Ag/Ag.
• Therefore, since Ecell is positive, EAg,Ag is
• placed before ECu2,Cu.

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Exp. 32 Introduction

• Standard Cell Potential (Eocell) is measured
• when Tsoln 25oC and ions 1 mol/L.
• The Nernst equation is applied to redox
• systems that are not at standard conditions
• Ecell Eocell (0.0592/n)logQ
• Where Q Cu2/Ag2 and n e-
• Cu(s) 2Ag(aq) ? 2Ag(s) Cu2(aq)

13
Exp. 32 Procedural Notes

• Part A.1. Set up as in Fig. 32.3. Use 50 mL
• beakers (obtain from stockroom).
• I have the potentiometers.
• Please try to keep the alligator clips out of the
  
• solutions.
• Part C. Omit

14
Exp. 32 Report Sheet

• Questions 1, 2, 4