Electrochemistry- study of interchange of chemical and electrical energy

1
Electrochemistry-study of interchange of
chemical and electrical energy

• Generating electric current from spontaneous
  chemical reactions and use of current to produce
  chemical change
• Labs
• 31 The Thermodynamics of the Dissolution of
  Borax
• 33 Electrolytic Cells Avogadros Number
• Chemical Equations
• Chapter 12

2
Oxidation-Reduction Reactions (Redox Reactions)

• Oxidation (LEO)
• Loss of electrons-atoms or ions undergo increase
  in oxidation state (is oxidized)
• Electrons given to another atom which is being
  reduced (reducing agent or reductant)
• Reduction (goes GER)
• Gain of electrons-atoms/ions undergo decrease in
  oxidation state (is reduced)
• Takes electrons away from another atom which is
  being oxidized (oxidizing agent or oxidant)

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• Redox reaction can be written as two
  half-reactions (one reduction, one oxidation)
• To generate current
• Separate oxidizing agent from reducing agent
• Electron transfer occurs through wire
• Electron transfer directed through device to
  provide useful work
• Reactants separated by salt bridge/porous
  partition
• Each half-reaction has a potential, or voltage,
  associated with it
• Given as reduction half-reactions
• Read in reverse and change sign on voltage to get
  oxidation potentials

4
Electrochemical cellsdevice associated with
redox reaction (galvanic cells, electrolytic
cells)

• Electrochemical process involves electron
  transfer at interface between electrode and
  solution
• Species undergoing reduction receive electrons
  from cathode
• Species in solution act as oxidizing agent
• Species undergoing oxidation donate electrons to
  anode
• Species in solution act as reducing agent

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mations/anodereaction.html
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mations/cathodereaction.html
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Galvanic Cell (voltaic cell)

• Chemical energy ? electrical energy
• Harnesses energy of spontaneous redox reactions
• Physically separate chemicals in 2 half-reactions
• Electrons generated by oxidation half-reaction
  flow through electrical conductor before being
  used in reduction half-reaction
• Flow diverted through meters, motors, light bulbs
  to perform useful work before reaching
  destination
• Current (defined by physicists as flow of
  positive charge) always in opposite direction
  from flow of electrons (always from anode to
  cathode)

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• Electrodes are metal strips
• Sign of electrodes determined by
• Since electrons flow out of anode and into
  external circuit, anode is negative
• Since electrons flow from external circuit into
  cathode, cathode is positive
• Opposite is true for electrolytic cells
• Where reactions occur (The Red Cat ate An Ox)
• Oxidation occurs at anode (AN OX)-on left
• Reduction occurs at cathode (RED CAT)-on right

7
Porous barrier separate two compartments

• Allows for migration of positive/negative ions
  between half-cells, completing electric circuit
• Problem with porous barriers
• Inside barriers, ionic solutions mix
• Has effect on operation of cell

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8
Galvanic cells with salt bridges

• Inverted tube contains electrolyte
• Gel (agar) added to provides
  firmness but permits ion flow
• Prevents two reacting solutions from mixing
• Ions dont react with other ions in cell or with
  electrode material
• Maintains electrical neutrality in system
• Provides - ions to equal ions created at anode
  (during oxidation) and ions to
  replace - ions being used up at
  cathode (during reduction)
• Anions always migrate toward anode
• Cations toward cathode

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Electron Flow
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Cell Potential (Ecell) or Electromotive force
(emf)

• Force with which electrons flow from - electrode
  (anode on left) to electrode (cathode on right)
  through external wire
• Due to PE difference of electrons before/after
  transfer
• In electrochemical cell, electric potential
  created between two dissimilar metals
• Greater tendency or potential of two
  half-reactions to occur spontaneously, greater
  emf of cell
• Measured in volts (V-why called cell voltage)
• 1 V 1 J/coulomb (of charge transferred)
• Measured with voltmeter which draws current
  through known resistance (heat is produced)

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• Voltage of voltaic cells
• All based on spontaneous chemical reactions
• ?G must always be negative
• Voltage of voltaic cell is always positive (EMF)
• Subtract smaller reduction potential from larger
  one
• Same as EMF cathode anode
• Under standard conditions, voltage of cell is
  same as total voltage of redox reaction
• Standard emf of standard cell potential (E0cell)
• Under nonstandard conditions, cell voltage
  computed by using Nernst equation
• As galvanic cell operates
• Redox reaction of cell approaches equilibrium
• Capacity to deliver useful electrical energy
  decreases
• At equilibrium, cell ceases to function (dead
  battery)

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Standard Reduction Potentials

• Cell potentials can be measured
• Half-cell potential cannot

15
Measuring Potential

• Galvanic cell under standard conditions made
  using arbitrary standard hydrogen electrode (SHE)
  and test half-cell w/different half reaction
• Assigned standard electrode potential of exactly
  0.00 V
• Half reaction always written as reduction
• Eo values corresponding to reduction
  half-reactions with all solutes standard
  reduction potentials-Eocell ?

Under ideal conditions where ideal behavior is
assumed
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Standard Reduction Potentials

• All solutions are 1M, gases at 1 atm, T 25oC
• Write oxidation/reduction half-reactions for cell
• Look of reduction potential (Eoreduction) for
  reduction half-reaction in table
• Half reaction w/higher reduction potential
• Look of reduction potential for reverse of
  oxidation half-reaction and reverse sign
  (Eooxidation -Eoreduction)
• Half reaction w/lower reduction potential/sign
  reversed
• Add potentials to get overall standard cell
  potential
• Two half reactions are balanced for electrons
  exchanged but value of each Eo remains unchanged
  (intensive property-does not depend on how many
  times reaction occurs)
• If sum positive, reaction is spontaneous/runs on
  own (always positive for electrochemical cells)
• If sum negative, energy supplied to make reaction
  go

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• Each half-reaction associated w/signed numerical
  value
• More positive it is, greater oxidizing power of
  redox half-reaction
• More negative it is, greater reducing power of
  reverse redox half-reaction
• Larger difference between Ered values, larger
  Ecell

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(1 atm)
19
Line Notation

• Anode written first on left/cathode on right
• Reactants written 1st on each side
• Vertical bar is boundary between two phases
• If both substances in same phase, separated by
  comma, not vertical bar
• Double line represents salt bridge or porous disk
• When platinum electrode present, placed at left
  and/or right end of cell diagram

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Using the table of standard reductions provided
write the equation for the reaction between the
following two half cells, and determine its
voltage

• Au (s) Au 3 (aq) Cu 2 (aq) Cu (s)
• Gold is higher on table and written as found on
  table
• Au 3 (aq) 3 e - ?Au (s) Eo 1.50 V
• Copper is lower so it's written as oxidation
  (sign reversed)
• Cu (s) ? Cu 2 (aq) 2 e- Eo -0.34 V
• Equation balanced for exchange of electrons (each
  needs 6)
• 2 Au 3 (aq) 6e - ? 2 Au (s) E 1.50 V
• 3 Cu (s) ?3 Cu 2 (aq) 6e E -0.34 V
• Size of values for Eo of reactions not
  changed
• Add two half reactions and E0 values
• 2 Au 3 (aq) 3 Cu (s) ? 2 Au (s) 3 Cu 2 (aq)
  
• E 1.50 V ( - 0.34 V ) 1.16 V

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Place the following in order of increasing
strength as oxidizing agents
-0.44 0.954 2.87 0.22
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Describe a galvanic cell based on the two
half-reactions below
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Homework

• Read 17.1-17.2, pp.827-837
• Q pp. 867-869, 14 a/b/c/f/i/k, 16 a/b/d/f, 26a,
  28 (a only), 30b, 32 (b only), 34a, 36b

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Cell Potential, Electrical Work, and Free Energy
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Complete Description of a Galvanic Cell

• Electrical energy delivered by galvanic cell
  equal to quantity of useful work obtained as
  result of cell operation
• Work, w, measured in relation to amount of
  charge, q, transferred between anode/cathode of
  cell
• This quantity, potential difference, E, is
  defined as E w/q.
• SI unit is joule per coulomb or volt (V)
• Cell potential (always positive for galvanic
  cell)
• Direction of electron flow (direction that yields
  potential)
• Designation of anode/cathode
• Nature of each electrode/ions present in
  compartments
• Chemically inert conductor (such as Pt) required
  if only ions are present (no substance in
  reaction is conducting solid)

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Electric work and cell potential

• Free energy change occurring during chemical
  reaction is measure of maximum work that system
  can perform
• Potential (E) -Work (w) / Charge (q)
• So w -qE
• Work leaving system has negative charge
• Faraday (F) charge in coulombs per mole of
  electrons (96,485 C/mol e-)
• Then q nF and w -nFE
• n number of moles of electrons

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• ?G (Gibbs free energy) is measure of spontaneity
  of process occurring at constant T/P
• emf, E, of redox reaction also indicates whether
  reaction is spontaneous
• From thermodynamics we know that
• ?G ?U T?S ?(PV) and U heat w
• Therefore, at constant T/P ?G w
• Therefore ?G -nFE and at standard state ?G0
  -nFE0 (relationship between emf/free energy
  changes )
• ?Go Standard Gibbs free energy change (kJ/mol
  or Joules)
• n moles of electrons exchanged in reaction
  (mol)
• F Faradays constant, 96,485 coulombs/mole (1
  mole of electrons has a charge of 96,485
  coulombs)
• Eo Standard reaction potential (V or
  Joules/Coulomb)

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• Since both n/F are
• value of E leads to value of ?G, which
  indicates spontaneous reaction
• If ?G and E have opposite signs, E predicts
  direction of reaction
• If Eo is positive, ?Go is negative (lt 0)-reaction
  spontaneous (has positive cell potentials)
• If Eo is negative, ?Go is positive (gt 0)-reaction
  is nonspontaneous (but is spontaneous in reverse
  direction)

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Relationship between thermodynamics (push behind
electrons) and electrochemistry

• Relationship between reaction potential and free
  energy for a redox reaction is given by
• Emf potential difference (V) work (J)
• charge (C)
• Driving force (emf) is defined in terms of
  potential difference (in V) between two points in
  circuit
• One coulomb is amount of charge that moves past
  any given point in circuit when current of 1
  ampere (amp) is supplied for one second (1 ampere
  1 coulomb/sec)
• Faradays law states that during electrolysis,
  passage of 1 faraday through circuit brings about
  oxidation of one equivalent weight of substance
  at one electrode (anode) and reduction of one
  equivalent weight at other electrode (cathode)

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• emf is not converted to work with 100 efficiency
• Energy always lost as heat, but wmax useful for
  calculating efficiency of conversion
• wmax -qEmax
• Relationship to free energy (energy driving
  reaction due to movement of charged particles
  giving rise to potential difference)
• wmax EG
• EG -qEmax -nFEmax
• EG -qEmax
• EG0 -nFE0
• When Ecell positive (spontaneous), EG will be
  negative (spontaneous), so there is agreement
• Standard cell potential, Eocell, measured and
  standard electrode potential of test half-cell
  determined by using Eocell Eocathode - Eoanode
• Eocathode-standard reduction potential for
  reaction occurring at cathode, represents
  tendency to remove es from electrode surface
• Eoanode-standard reduction potential for reaction
  occurring at anode and represents its tendency to
  remove es from anode

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Dependence of Cell Potential on Concentration

• Cell voltages at nonstandard concentrations (not
  1 M)

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Nernst Equation-way to relate E0 at standard
conditions and E, potential at any real condition

• Standard state impossible to achieve in reality
• As soon as wire hooked to two half-cells,
  reaction proceeds and changes concentrations of
  all reactants and products
• Heating or cooling makes reaction deviate from
  standard temperature

37
Calculating Nernst Equation

• From thermodynamics, recall G Go RTlnQ
• If we divide everything by nF
• since ?G nFE, or E ?G/(nF)
• Ecell cell potential at non-standard conditions
• E0cell standard reduction potential
• R 8.314 J/molK (the gas constant)
• F 96485 coul/mol (Faraday's constant)
• T absolute temperature
• n number of moles of electrons transferred in
  balanced equation
• Q reaction quotient for reaction aA bB ? cC
  dD
• Expressed in terms of base 10 rather than ln
  (standard conditions of 298K)
• Can be used to find cell potential at any set of
  conditions
• Cells spontaneously discharge until they achieve
  equilibrium (at equilibrium, cell is dead)

38
Consider the Daniell Cell at 25 CZn(s)
Cu2(aq) Cu(s) Zn2(aq)

• Find cell potential at following conditions when
  Cu2 1.00 M, Zn2 1.0109 M and when
  Cu2 0.10 M, Zn2 0.90 M.
• Recall that standard potential for Daniell cell
  is Eo 1.10 V
• Nernst equation used to find potentials at
  nonstandard conditions

When Cu2 1.00 M, Zn2 1.0109 M
When Cu2 0.10 M, Zn2 0.90 M
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Depiction of concentration cell
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• Cell in which current flows due only to
  difference in concentration of ion in 2 different
  compartments of cell
• Le Châteliers principle used to determine effect
  on potential
• Shift to left reduces potential
• Shift to right increases potential
• If concentrations are different, stress is put on
  system that will be equalized by electron flow to
  allow reduction and oxidation to occur
• Voltages typically small
• When concentrations in half-cells become equal,
  E0cell 0 and system is at equilibrium

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46
Calculate the EMF of the cell Zn(s) Zn2 (0.024
M) Zn2 (2.4 M) Zn(s)

• Zn2 (2.4 M) 2 e Zn Reduction
• Zn Zn2 (0.024 M) 2 e Oxidation
• Zn2 (2.4 M) Zn2 (0.024 M), DE 0.00
• Using Nernst equation
• (0.024) DE 0.00 - 0.0592/2 log (0.024/(2.4)
• (-0.296)(-2.0)
• 0.0592 V

47
Show that voltage of electric cell is unaffected
by multiplying reaction equation by positive
number

• Mg Mg2 Ag Ag
• Mg 2 Ag Mg2 2 Ag
• DE DEo 0.0592/2 log Mg2/Ag2
• 2 Mg 4 Ag 2 Mg2 4 Ag
• DE DEo 0.0592/4 log Mg22/Ag4
• Simplified to the 1st equation, showing cell
  potential DE not affected

48
Calculation of Equilibrium Constants for Redox
Reactions

• The standard reaction potential is related to the
  equilibrium constant
• At equilibrium, Ecell 0 and Q K
• As cells discharge, concentration changes, Ecell
  changes.
• For a cell at concentrations and conditions other
  than standard, a potential can be calculated
  using the Nernst equation
• If Eo is positive, then K gt 1 and forward
  reaction favored
• If Eo is negative, then K lt 1 and reverse
  reaction is favored

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50
The standard cell potential dE for the reaction
Fe Zn2 Zn Fe2 is -0.353 V. If a piece of
iron is placed in a 1 M Zn2 solution, what is
the equilibrium concentration of Fe2?

• Equilibrium constant K may be calculated using K
  10(n DE)/0.0592
• 10-11.93
• 1.2x10-12
• Fe2/Zn2. Since Zn2 1 M, it is
  evident that Fe2 1.2-12 M

51
Homework

• Read 17.3-17.4, pp. 837-846
• Q pp. 869-871, 38, 40, 46, 48, 54, 60, 66, 70

52
Batteries

• Portable, self-contained electrochemical power
  source (DC) consisting of one or more voltaic
  cells, connected in series
• Greater voltages achieved by using multiple
  voltaic cells in single battery (12V)
• When connected in series, battery produces
  voltage that is sum of emfs of individual cells
• Higher emfs achieved by using multiple batteries
  in series
• Electrodes marked (cathode) and (anode)

53
Lead-acid storage battery

• As battery discharged, uses up sulfate
  ions/electrodes become coated w/lead sulfate
• Reverse reactions regenerate sulfate ion in
  solution/reduce amount of lead sulfate
  contaminating electrode surfaces
• Each pair produces 2 volts (6 pairs of
  electrodes used in 12-volt car battery)
• When jump starting car, connect ground cable on
  dead car to metallic contact away from battery.
  Otherwise, could explode
• Causes electrolysis of water/production of H2/O2
  which could ignite

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Common Dry Cell Battery(acid version)

• Zinc-anode
• Carbon rod in contact with moist paste-cathode

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Alkaline dry cell

• NH4Cl replaced w/KOH or NaOH
• Last longer than acid cells because zinc (anode)
  corrodes more slowly in basic environment
• Cathode (graphite rod) inserted into paste made
  of manganese dioxide, water and potassium
  hydroxide
• Zn(s) 2OH-(aq) ? ZnO(s) H2O 2e- (anode)
• 2MnO2(s) H2O 2e- ? Mn2O3(s) 2OH-(aq)
  (cathode)
• Total voltage is 1.54 volts
• Not rechargeable

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Mercury Battery
58

• Fuel Cells
• Galvanic cells where reactants continuously
  supplied
• Energy normally lost as heat is captured and used
  to produce an electric current
• Redox reaction
• Hydrogen oxidized at anode and oxygen reduced at
  cathode to form water and electricity
• 2x as effective as gas, oil or coal-powered
  generators in converting chemical energy into
  electricity

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Corrosionoxidation of metal

• Oxidation of most metals by oxygen is spontaneous
  redox reactions
• Many metals develop thin coating of metal oxide
  on outside that prevents further oxidation
• Some metals, such as copper, gold, silver and
  platinum (noble metals), are relatively difficult
  to oxidize

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Corrosion of Iron

• Anodic regions
• Regions of steel alloy where iron is more easily
  oxidized
• Fe ? Fe2 2e-
• Cathodic regions
• Areas resistant to oxidation
• Electrons flow from anodic regions react
  w/oxygen
• O2 2H2O 4e- ? ?4OH-
• Presence of water essential to iron corrosion
• Presence of salt accelerates corrosion by
  increasing electron conduction from anodic to
  cathodic regions

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Prevention of Corrosion

• Coating w/metal that forms oxide coat to protect
  metal that would not develop protective coat
• Galvanizing
• Place sacrificial of more easily oxidized metal
  on top of metal to protect
• Zinc over iron
• Alloying
• Addition of metals that change steels reduction
  potential.
• Nickel and chromium alloyed to iron

64

• Cathodic Protection
• Connection of easily oxidized metals (an anode)
  to less easily oxidized metals keeps less from
  experiencing corrosion
• Anode corrodes-must be replaced periodically
• Magnesium as anode to iron pipe
• Titanium as anode to steel ships hull

65
Electrolysis

• Decomposition of substance by electric current

66
Electrolytic cells
http//www.infoplease.com/chemistry/simlab/electro
lpt3.html

• Nonspontaneous reactions
• Electrical energy required to induce reaction
• Two electrodes immersed in electrically
  conductive sample
• Electrical voltage (gt1.10 V) applied to them
• Voltage increased until electrons flow in
  opposite direction (electrolytic)
• At cathode-reduction occurs (RED CAT)
• At anode-oxidation occurs (AN OX)
• Electrical energy is converted into chemical
  energy
• Electrolytic cells are used for electroplating

67

• (a) Standard galvanic cell based on spontaneous
  reaction
• Zn Cu2 ? Zn2 Cu
• (b) Standard electrolytic cell Power source
  forces opposite reaction
• Cu Zn2 ? Cu2 Zn

68
What voltage is necessary to force the following
electrolysis reaction to occur?

• 2I-(aq) Cu2(aq)? I2(s) Cu(s)
• Which process would occur at the anode? Cathode?
  Assuming the iodine oxidation takes place at a
  platinum electrode, what is the direction of
  electron flow in this cell?

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Calculating Quantity of Substance Produced or
Consumed

• To determine quantity of substance either
  produced or consumed during electrolysis given
  time known current flowed
• Write balanced half-reactions involved
• Calculate number of moles of electrons that were
  transferred
• Calculate number of moles of substance that was
  produced/consumed at electrode
• Convert moles of substance to desired units of
  measure

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A 40.0 amp current flowed through molten
iron(III) chloride for 10.0 hours (36,000 s). 
Determine the mass of iron and the volume of
chlorine gas (measured at 25oC and 1 atm) that is
produced during this time.

• Write half-reactions that take place at
  anode/cathode
• anode (oxidation)  2 Cl-  Cl2(g) 2 e
• cathode (reduction)  Fe3 3 e-   Fe(s)
• Calculate number of moles of electrons
• Calculate moles of iron/chlorine produced using
  number of moles of electrons calculated and
  stoichiometries from balanced half-reactions.  (3
  moles electrons produce 1 mole of Fe/2 moles of
  electrons produce 1 mole of chlorine gas)
• Calculate mass of iron using molar mass and
  calculate volume of chlorine gas using ideal gas
  law (PV nRT).

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Calculating Time Required

• To determine quantity of time required to produce
  known quantity of substance given amount of
  current that flowed
• Find quantity of substance produced/consumed in
  moles
• Write balanced half-reaction involved
• Calculate number of moles of electrons required
• Convert moles of electrons into coulombs
• Calculate time required

73
How long must a 20.0 amp current flow through a
solution of ZnSO4 in order to produce 25.00 g of
Zn metal

• Convert mass of Zn produced into moles using
  molar mass of Zn
• Write the half-reaction for the production of Zn
  at the cathode Zn2(aq) 2 e-  Zn(s)
• Calculate moles of e- required to produce moles
  of Zn using stoichiometry of the balanced
  half-reaction (2 moles of electrons produce 1
  mole of zinc)
• Convert moles of electrons into coulombs of
  charge using Faraday's constant
• Calculate time using current and coulombs of
  charge

74
Calculating Current Required

• To determine amount of current necessary to
  produce known quantity of substance in given
  amount of time
• Find quantity of substance produced/or consumed
  in moles
• Write equation for half-reaction taking place
• Calculate number of moles of electrons required
• Convert moles of electrons into coulombs of
  charge
• Calculate current required

75
What current is required to produce 400.0 L of
hydrogen gas, measured at STP, from the
electrolysis of water in 1 hour (3600 s)?

• Calculate number of moles of H2
• Write equation for half-reaction that takes
  place. Hydrogen produced during reduction of
  water at cathode.  Equation for this
  half-reaction is 4 e- 4 H2O(l)  2 H2(g) 4
  OH-(aq)
• Calculate number of moles of electrons (4 mole of
  e- required to produce 2 moles of hydrogen gas,
  or 2 moles of e-'s for every one mole of hydrogen
  gas)
• Convert moles of electrons into coulombs of
  charge
• Calculate current required

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How many grams of copper can be reduced by
applying a 3.00 A current for 16.2 min to a
solution containing Cu2 ions?

• Time ? current ? Coulombs ? moles e- ? moles Cu ?
  g Cu
• 16.2 min 60 sec 3 C 1 mol e- 1 mol Cu
  63.54 g Cu
• 1 min 1 sec 96,486 C 2
  mol e- 1 mol Cu
• 0.96 g Cu

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Electrolysis can be used to separate mixture of
ions, if reduction potentials are fairly far apart

• Remember, metal ion with highest reduction
  potential is easiest to reduce
• Predict order of reduction and which of following
  ions will reduce first at cathode of electrolytic
  cell Ag, Zn2, IO3-

78
Electroplating

• Deposit neutral metal atoms on electrode by
  reducing metal ions in solution
• One metal coated with another
• Presence of active electrode that takes part in
  electrolysis reaction
• Anode-piece of plating metal
• Cathode-object to be plated
• Plating solution is NiSO4 because SO42- ion does
  not participate in plating reaction

79
How long must a current of 5.00 A be applied to a
solution of Ag to produce 10.5 g silver metal?

• 10.5 g Ag 1 mol Ag 1 mol e- 96,485 C 1
  sec 1 min
• 107.868 g Ag 1 mol Ag 1 mol e-
  5 C 60 sec
• 31.3 min

80
Electrolysis of Water
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mations/electrolysisofwater.html

• Requires soluble salt/dilute acid to serve as
  electrolyte
• Anode
• 2H2O(l) ? O2(g) 4H(aq) 4e- Eoox -1.23 V
• Cathode
• 2H2O(l) ? H2(g) 2OH-(aq) Eored -0.83 V
• Overall reaction
• 6H2O(l) ? 2 H2(g) O2(g) 4H(aq) 4e-
  Eocell -2.06 V
• If in single container, H/OH- combine to yield 4
  additional water molecules

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What volume of H2(g) and O2(g) is produced by
electrolyzing water at a current of 4.00 A for
12.0 minutes (assuming ideal conditions)?

• 2H2O(l) ? ?2H2(g) O2(g)
• Actual ratio is not exactly 21 for a variety of
  reasons including oxygen solubility.
• 12 min 60 sec 4 C 1 mol e- 1 mol H2
  22.4 L
• 1 min 1 sec 96,486 C 2 mol e-
  1 mol H2
• 0.334 L H2 ? 0.167 L O2

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Electrolysis of molten salts

• NaCl
• Good conductor as liquid
• Melting salt frees ions
• Makes it electrically conductive
• Ions of opposite charge migrate to these
  electrodes and react

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Electrolysis of aqueous solutions

• Aqueous solutions of salts are electrically
  conductive and can be electrolyzed
• For solutions, two possible reactions occur
  (water can be both oxidized and reduced)
• At cathode
• If metal ion is very active metal, water will be
  reduced (2H2O 2e- ? H2 2OH-)
• If metal ion is inactive or active metal, metal
  ion will be reduced
• At anode
• Oxidation of salts anion or (?)
• Oxidation of water (2H2O ? O2 4H 4e-)

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• To determine which will occur at anode
• If anion is polyatomic ion, it generally will not
  be oxidized
• SO42-/NO3-/ClO4- not oxidized in aqueous solution
• Cl-/Br-/l- will be oxidized in aqueous solution
• If anion in one salt is oxidized in aqueous
  electrolysis, that same anion in any other salt
  will also be oxidized
• If solution of NaBr results in Br- being oxidized
  to Br2, predict that solutions of KBr, CaBr2,
  NH4Br and AlBr3 will all produce Br2 at anode

85
Commercial Electrolytic Processes

• Abundance of elements on earth
• 1st Oxygen
• 2nd Silicon
• 3rd Aluminum (very active metal so difficult and
  expensive originally to purify)

• Since metals are easily oxidized, most found as
  ores, mixtures of ionic compounds. Au, Ag, and
  Pt are more difficult to oxidize, so often found
  as pure metals.

86
Production of aluminum from molten-salt
electrolysis (purification of aluminum from
bauxite ore)
Hall-Heroult process
87

• Electrorefining (purifying) metals
• Copper ore is refined by roasting
• Impure copper is anode
• Small strip pure copper is cathode
• During electrolysis, copper is oxidized to Cu2
  at anode and then reduced to copper metal again
  at cathode
• Impurities such as silver and gold drop to bottom
  as sludge which is then salvaged

88

• Metal Plating
• Electroplating thin layers of decorative metal on
  less expensive metal (silver and gold onto iron,
  chromium on to car parts for decoration and
  resistance to corrosion)

89
Electrolysis of concentrated aqueous sodium
chloride solutions (brine) produces hydrogen and
hydroxide ions at cathode and chlorine gas at
anode

• If electrodes separated by porous membrane, H2,
  NaOH, and Cl2 produced
• If solution stirred, chlorine gas reacts with
  sodium hydroxide to form sodium hypochlorite
  (NaOCl) solution (bleach)
• Electrolysis of molten NaCl produces sodium metal
  and chlorine gas (Downs cell)

90
Homework

• Read 17.5-17.8, pp. 846-866
• Q pp. 871-872, 74, 76, 78, 80, 84, 88 a (dont
  forget water)
• Do 1 additional exercise and 1 challenge problem
• Submit quizzes by email to me
• http//www.cengage.com/chemistry/book_content/0547
  125321_zumdahl/ace/launch_ace.html?folder_path/ch
  emistry/book_content/0547125321_zumdahl/acelayer
  actsrcch17_ace1.xml
• http//www.cengage.com/chemistry/book_content/0547
  125321_zumdahl/ace/launch_ace.html?folder_path/ch
  emistry/book_content/0547125321_zumdahl/acelayer
  actsrcch17_ace2.xml
• http//www.cengage.com/chemistry/book_content/0547
  125321_zumdahl/ace/launch_ace.html?folder_path/ch
  emistry/book_content/0547125321_zumdahl/acelayer
  actsrcch17_ace3.xml