Electrochemistry

1
Electrochemistry

• Electrochemistry - study of the relationships
  which exist between the flow of electrons and
  chemical reactions
• Types of electrochemical systems
• electrolytic - chemical reaction which occurs
  when electrical current is passed through
  solution
• voltaic/galvanic - spontaneous reactions able to
  generate a supply of electricity (e.g.,
  batteries)
• Spontaneous redox reactions (see pp. 126-131) are
  coupled in such a way (i.e., an electrochemical
  cell) as to allow electrons to flow through an
  external circuit
• The electrochemical cell design half-cells (2)
  salt bridge potentiometer electrodes
  electrolyte solutions conducting wire

2
Electrochemistry

• An electrochemical cell is a system consisting of
  electrodes that dip into an electrolyte in which
  a chemical reaction either uses or generates an
  electric current.
• A voltaic, or galvanic, cell is an
  electrochemical cell in which a spontaneous
  reaction generates an electric current.
• An electrolytic cell is an electrochemical cell
  in which an electric current drives an otherwise
  nonspontaneous reaction.

3
Electrochemistry

• A complete redox reaction takes place in a
  galvanic cell
• Overall reaction separated into half-reactions
  which take place at the anode and cathode
• Given the following reaction
• Zn(s) Cu2(aq) ? Zn2(aq) Cu(s)
• the two half-reactions are
• oxidation half-reaction Zn(s) ? Zn2(aq)
  2e- (anode)
• reduction half-reaction Cu2(aq) 2e- ? Cu(s)
  (cathode)
• Electrode reactions
• Anode site of oxidation electrons originate
  there neg. pole of cell (anions migrate toward)
• Cathode site of reduction electrons consumed
  there pos. pole of cell (cations migrate toward)

4
Figure 20.2 Atomic view of voltaic cell. In a
voltaic cell, two half-cells are connected in
such a way that electrons flow from one metal
electrode to the other through an external
circuit.
The electrons flow through the external circuit
to the copper electrode where copper ions gain
the electrons to become copper metal.
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Figure 20.3 Two electrodes are connected by an
external circuit
Daniel Cell Design
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Cell Notation

• Cell notation is used to describe structure of
  galvanic cell
• For the Zn/Cu cell, the galvanic cell notation
  is
• Zn(s) ?Zn2(aq) Cu2(aq) ?Cu(s)
• ? phase boundary
• salt bridge
• anode reaction to the left of the salt bridge
• cathode reaction to the right of the salt
  bridge
• both half-cell reactions in order of spontaneous
  reaction
• Zinc solid reacts to form zinc(II) ion at the
  anode
• Copper(II) ion reacts to form copper metal at
  the cathode

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Quiz

• Which of the following statements is incorrect
• In an electrolytic cell, reduction occurs at the
  anode.
• Aluminum metal would form at the cathode during
  the electrolysis of molten AlBr3.
• The cathode is labeled "" in a voltaic cell.
• Oxidation occurs at the anode in a voltaic cell.
• Electrons flow from the anode to the cathode in
  all electrochemical cells.

8
Quiz

• Consider the following notation for an
  electrochemical cell
• ZnZn2 (1M)Fe3 (1M), Fe2 (1M)Pt
• What is the balanced equation for the cell
  reaction?
• Zn(s) 2Fe3(aq) ? 2Fe2(aq) Zn2(aq)
• Zn2(aq) 2Fe2(aq) ? Zn(s) 2Fe3(aq)
• Zn(s) 2Fe2(aq) ? 2Fe3(aq) Zn2(aq)
• Zn(s) Fe3(aq) ? Fe2(aq) Zn2(aq)
• Zn(s) Fe2(aq) ? Fe(s) Zn2(aq)

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Electromotive Force

• The movement of electrons is analogous to the
  pumping of water from one point to another.

• Water moves from a point of high pressure to a
  point of lower pressure. Thus, a pressure
  difference is required.

• The work expended in moving the water through a
  pipe depends on the volume of water and the
  pressure difference.

10
Electromotive Force

• The movement of electrons is analogous to the
  pumping of water from one point to another.

• An electric charge moves from a point of high
  electrical potential (high electrical pressure)
  to one of lower electrical potential.

• The work expended in moving the electrical charge
  through a conductor depends on the amount of
  charge and the potential difference.

11
Electromotive Force

• Potential difference is the difference in
  electric potential (electrical pressure) between
  two points.

• You measure this quantity in volts.

12
Electromotive Force

• The Faraday constant, F, is the magnitude of
  charge on one mole of electrons it equals 96,500
  coulombs (9.65 x 104 C).

• In moving 1 mol of electrons through a circuit,
  the numerical value of the work done by a voltaic
  cell is the product of the Faraday constant (F)
  times the potential difference between the
  electrodes.

13
Electromotive Force

• The Faraday constant, F, is the magnitude of
  charge on one mole of electrons it equals 96,500
  coulombs (9.65 x 104 C).

• In the normal operation of a voltaic cell, the
  potential difference (voltage) across the
  electrodes is less than than the maximum possible
  voltage of the cell.

• The actual flow of electrons reduces the
  electrical pressure.

14
Electromotive Force

• The Faraday constant, F, is the magnitude of
  charge on one mole of electrons it equals 96,500
  coulombs (9.65 x 104 C).

• In the normal operation of a voltaic cell, the
  potential difference (voltage) across the
  electrodes is less than than the maximum possible
  voltage of the cell.

• Thus, a cell voltage has its maximum value when
  no current flows.

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Electromotive Force

• The maximum potential difference between the
  electrodes of a voltaic cell is referred to as
  the electromotive force (emf) of the cell, or
  Ecell.

• It can be measured by an electronic digital
  voltmeter which draws negligible current.

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Electromotive Force

• Electrons are driven (pushed) through
  conducting wire in the direction of anode ?
  cathode by cell force
• Origin of cell force is maximum electric
  potential difference between electrodes or
  electromotive force (Ecell) or cell potential
• Potential difference - difference in electrical
  potential (electrical pressure) between two
  electrodes standard unit of cell potential
  difference is the Volt
• Electrical work
• electrical work charge moved X potential
  difference
• J C X V
• wmax ?G -nFEcell

17
Standard Cell Potentials

• Standard potential of galvanic cell sum of
  standard half-cell potentials of ox at anode and
  red at cathode
• Eocell Eoox(anode) Eored(cathode)
• Since -Eored(anode) Eoox(anode)
• Eocell Eored(cathode) - Eored(anode)
• For a spontaneous cell reaction, Eocell is
  positive (since Gibbs free energy change must be
  lt 0)
• Method developed to estimate standard cell
  potentials under standard conditions (1 atm, 1 M,
  25 oC)
• Standard cell potentials termed standard
  reduction potentials according to above formula
• SHE used to determine standard reduction
  potentials

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Figure 20.5 Hydrogen Electrode
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Standard Cell Potentials

• Ecell driving force for cell reaction
  (pressure)
• Ecell is a function of cell reaction and
  electrolyte concs.
• Standard cell, Eo - emf of a galvanic (voltaic)
  cell operating under standard state conditions
• SHE - standard hydrogen electrode by convention
  Eo for SHE 0
• Coupled reactions with SHE provide standard
  reduction potentials
• Eored 0 - Eo1/2 cell (SHE as cathode)
• Eored Eo1/2 cell - 0 (SHE as anode)
• Strongest oxidizing agent - largest Eo value
  (most -)
• Strongest reducing agent - smallest Eo value
  (most )

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21
Quiz

• Consider the electrochemical cell
• Zn(s) Zn2(aq) Br-(aq) Br2(l) Pt
• Calculate the standard potential of this cell,
  Ecell.
• Eo (Zn2/Zn) -0.76 Eo (Br2/Br-) 1.07
  
• 0.31 V
• 1.83 V
• -0.31 V
• 1.30 V
• none of the above

22
Quiz

• Given the following reduction potentials
• What is the standard cell potential (Eocell) for
  the voltaic cell based on the reaction Fe(s) 2
  Fe3(aq) ? 3 Fe2(aq)
• 0.33 V b. 1.98 V c. 1.21 V d. -0.11 V
• e. -0.55 V

23
Quiz

• Given the following reduction potentials
• Which of the following is the strongest oxidizing
  agent?
• Sn2 b. Fe3 c. Fe2 d. Cr(s)
• e. Sn4

24
Quiz

• How much electrical work is performed through the
  consumption of 150 g of zinc in the following
  voltaic cell at standard conditions?
• Zn(s) Zn2(aq) Br-(aq) Br2(l) Pt
• Eo (Zn2/Zn) -0.76 Eo (Br2/Br-) 1.07
  
• 1,050 kJ
• 4,100 kJ
• 474 kJ
• 810 kJ
• none of the above

25
Quiz

• The standard cell potential (Eocell) for the
  reaction below is 0.126 V. The value of ?Go for
  the reaction
• is _____ kJ/mol.
• Pb(s) 2 H(aq) ? Pb2(aq) H2(g)
• -24
• 24
• -12
• 12
• -50

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Cell Potentials Nernst

• Cell potentials are a function of half-cell
  electrolyte concentrations
• ?G RT ln Q/K
• And ?G -nFEcell
• Thus, Ecell -?G/nF
• Ecell -RT/nF ln Q/K
• Ecell RT/nF ln Q RT/nF ln K ( Eocell)
• Ecell Eocell RT/nF ln Q (Nernst Equation)

27
Quiz

• Consider the following electrode potentials
• Mg2 2e ? Mg E 2.37 V
• V2 2e ? V E 1.18 V
• Cu2 e ? Cu E 0.15 V
• Which one of the reactions below will proceed
  spontaneously from left to right?
• a. Mg2 V ? V2 Mg
• b. Mg2 2Cu ? 2Cu2 Mg
• c. V2 2Cu ? V 2 Cu2
• d. V 2Cu2 ? V2 2Cu
• e. none of these

28
Quiz

• At 25C, calculate the voltage of the cell if
  Ecell 0.460 V.
• Cu?Cu2(0.10 M) Ag(0.10 M) ?Ag(s)
• a. 0.282 V
• b. 0.371 V
• c. 0.430 V
• d. 0.460 V
• e. 0.490 V

29
Quiz

• What is the cell voltage (Ecell) for the
  following galvanic cell at 25 oC? (Eocell 0.15
  V)
• Cd(s)Cd2(0.026 M)Ni2(0.00420 M)Ni(s)
• a. 0.15 V
• b. 0.17 V
• c. 0.22 V
• d. 0.13 V
• e. 0.19 V

30
Quiz

• What is the pH of the test solution when Ecell
  0.612 V at 25 oC? Eo(AgCl/Ag) 0.22
• PtH2(g)(1 atm)H(test soln)AgCl(s),Ag(s)Cl-(2
  .80 M)
• a. 3.72
• b. 10.09
• c. 7.08
• d. 4.76
• e. 12.22

31
Quiz

• In an electrolytic cell, how many grams of Cu
  could be plated out of a CuSO4 solution at a
  current of 5.00 A for 2.00 min? (F 96500 C/mol)
• 318 g
• 0.395 g
• 0.329e-3 g
• 0.198 g
• 5.31 g

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Quiz

• How many minutes does it take to plate 0.800 g of
  silver metal onto a serving tray from an aqueous
  solution of AgNO3 at a current of 2.50 A? (F
  96500 C/mol)
• lt 2 minutes
• 2.38 minutes
• 4.77 minutes
• 9.54 minutes
• 23.8 minutes

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Figure 20.7 Relationships Among K, ? G, and
Ecell