Electrochemistry Review

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Electrochemistry Review

•  
• Goals
•  
• In this lab you will accomplish three things
  Build a standard voltaic (galvanic) cell and
  measure the potential.
• Determine the concentration of an unknown
  solution using the potential measured from a
  voltaic cell.
• Measure the voltage of a series of solutions with
  different hydronium ion concentrations and thus
  learn something about the relationship between
  the Nernst equation and the values obtained from
  a pH meter.

2
Introduction

• Electrochemistry deals with chemical reactions
  that produce electricity and chemical reactions
  that are propagated by an electrical current.
• In a voltaic cell electricity is produced from
  the chemical reaction
• In an electrolytic cell electricity is pumped
  into the cell to get the chemical reaction to
  proceed.
• used for electroplating metals such as
• chrome, copper or silver onto other surfaces
• as well as for the isolation of materials.

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Introduction

• The following equation shows the reaction for the
  isolation of sodium metal and chlorine gas from
  molten sodium chloride.
• To force the reaction to the right, current must
  be applied to this electrolytic cell.
• 2 NaCl (s) 2 Na o (s) Cl2 (g)

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Introduction

• In the experiment we will do in the laboratory,
  we will deal with only a voltaic cell.
• For every electrochemical reaction something must
  be oxidized and something must be reduced.
• The electrode where oxidation occurs is called
  the anode
• The electrode where reduction occurs in called
  the cathode
• Whether the anode is positive or negative depends
  upon whether the cell is electrolytic or voltaic.
  
• Since most of us are exposed to voltaic cells
  such as batteries in our everyday lives, it is
  common to think of the anode as negative but that
  is true ONLY in voltaic cells

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• Figure 1 is a schematic of a voltaic cell.

• In this voltaic cell the zinc metal is with the
  zinc ions and the copper metal is with the copper
  ions with a salt bridge connecting the two.
• Note the two half-cells.
• Each half of the cell is isolated from the other
  and the circuit is completed by a salt bridge
  connecting the two half cells.

• The salt bridge allows for the slow transfer of
  ions through a medium such as agar to which an
  electrolyte such as potassium nitrate has been
  added.
• This keeps each half-cell isolated in other
  words, the reduction reaction separated from the
  oxidation reaction.

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Line Notation

• For voltaic cells, it is common to use a
  shorthand notation to indicate what is going on
  in the cell.
• Below is given the shorthand for the voltaic cell
  in Figure 1.
• Zn Zn2 (1.0M) Cu 2 (1.0M) Cu

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Standard hydrogen electrode (SHE)

• By isolating a reaction in a half-cell it is
  possible to calculate what is called a standard
  reduction potential.
• The standard hydrogen electrode (SHE) is
  comprised of a platinum electrode immersed in 1.0
  M H solution with hydrogen gas at 1 atmosphere
  of pressure bubbled over the electrode.
• The SHE can act as an anode where H2 is oxidized
  to H, or it can act as a cathode where the H is
  reduced to H2.
• In either reaction the SHE is assigned a
  potential of ZERO volts.
• Standard reduction potentials are measured
  against the SHE with the corresponding electrode
  in a solution that is 1.0M.

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Table 1 shows some standard reduction potentials
(SRP)
Element Electrode Reaction Volts
Mg Mg2 2 e- ? Mg -2.37
Al Al3 3 e- ? Al -1.66
Zn Zn2 2 e- ? Zn -0.763
Fe Fe2 2 e- ? Fe -0.44
Ni Ni2 2 e- ? Ni -0.25
Sn Sn2 2 e- ? Sn -0.14
Pb Pb2 2 e- ? Pb -0.126
H2 2H 2 e- ? H2 0.000
Cu Cu2 2 e- ? Cu 0.337
I2 I2 2 e- ? 2 I - 0.535
Fe3 Fe 3 1 e- ? Fe 2 0.771
Ag Ag 1 e- ? Ag 0.799
Cl- Cl2 2 e- ? 2 Cl - 1.36

• As the reduction potential voltages increase so
  does the tendency for the ease of reduction of
  the species.

• Thus the reactions at the bottom of the table
  with high reduction potentials are strong
  oxidizing agents.

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What if the solutions are at some other
concentration or pressure?

• SRPs given for reactions are at 1.0 M for
  aqueous solutions and 1 atm for gases
• The Nernst equation (Equation 1) can be used to
  calculate SRPs at different concentrations and
  pressures.

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Nernst equation

• n is the number of electrons transferred
• F is Faradays constant
• Q is the reaction quotient.
• it is the ratio of the multiplicative product of
  the concentrations of the products to the
  multiplicative product of the concentrations of
  the reactants, each raised to the exponent of its
  stoichiometric coefficient.

• E is the observed potential under non-standard
  conditions
• E0 is the standard potential
• the value 2.303 converts the equation from
  natural logarithm (ln) to log10
• R is 8.314 J/molK
• T is temperature in Kelvin

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Nernst equation

• Assuming a temperature of 298K, Equation 1
  simplifies to Equation 2.

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Example Calculation

• Consider as an example a cell formed from the two
  half reactions Fe3 (1.00M)/ Fe2 (0.20M) and Cl2
  (5 atm)/Cl- (2.0 x 10-2M).
• First the standard cell potential, E0, can be
  calculated using the standard reduction
  potentials given in Table 1.
• A quick look at the table lets us know that for
  the standard potential of the cell to be positive
  (a requirement for a spontaneous reaction), the
  chlorine will be reduced to chloride ions and the
  iron II will be oxidized to iron III.
• Therefore the equation involving Fe2 and Fe3
  must be reversed as well as the sign of the
  potential given in the table.

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Example Calculation

• The two half reactions can be summed as below to
  determine E0.
• The factor of 2 must be used in the second
  equation so that no electrons are left in the net
  ionic equation, but the potential is not
  multiplied by the factor of 2 because potential
  is the ratio of energy to charge so any factor is
  canceled.
• Cl2 2e- ? 2Cl- 1.36 V
• 2Fe2 ? 2Fe3 2e- -0.771 V
• ______________________________________
• Cl2 2Fe2 ? 2Cl- 2Fe3 0.589 V

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What is Q?

• For this reaction,
• Q Cl-2 Fe32
• Cl2 Fe22

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Calculating the Cell Potential

• Using Equation 2 and substituting in the
  conditions for the nonstandard cell given in the
  example above, the potential for the cell can be
  calculated
•  
• E 0.589 - 0.0592 log10 2.0 x 10-22 1.002
  0.669 V 2
  5 0.202
•  

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Procedure I

• Construct a simple voltaic cell using 0.1 M
  copper sulfate CuSO4 solution and 0.1 M zinc
  sulfate ZnSO4 solution.
• You will not be using 1.0 M standard solutions as
  you might have expected from the discussion of
  SRP. The reason for this is that problems are
  encountered with the experiment at concentrations
  of this magnitude.
• You will use the Nernst equation to calculate the
  standard potential of 0.1 M cells.

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Procedure I

• Using the 0.1 M copper solution and a zinc
  solution of unknown concentration, construct
  another voltaic cell and record the potential as
  above.
• Use the Nernst equation to determine the
  concentration of the zinc solution.
• Using the 0.1 M zinc solution and a copper
  solution of unknown concentration, construct
  another voltaic cell and record the potential.
• Use the Nernst equation to determine the
  concentration of the copper solution.

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Procedure II

• Calibrate a pH meter using a two point
  calibration and two selected buffers (pH  4 pH
  7).
• Once you have calibrated the pH meter, switch the
  readings from pH to millivolts and record the
  millivolts for the series of buffers provided,
  including the pH 4 and pH 7.
• After you have collected the data, plot the
  voltage of each solution in volts as a function
  of log10H3O.
• Refer to the section on Concentration Cells at
  the end of section 21.4 in Silberberg if you need
  any help with the final question for this lab.