Chapter 21 Electrochemistry

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Chapter 21 Electrochemistry

• AP Chemistry
• Milam

2
Overview

• In physics electricity typically deals with
  electrons flowing through metals, the flow of
  charge is called current
• In chemistry we look at what causes electrons to
  flow, construction of batteries and which
  chemicals will work best for a specific function
  such as a battery
• Some concepts from both will be present and some
  unique to chemistry will be there as well

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Overview

• An electrolytic cell is one where external energy
  is used to cause a nonspontaneous chemical
  reaction to occur
• A voltaic cell is where a spontaneous chemical
  reaction (usually redox) produces electricity and
  can supply it to something external

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21-1 Electrical Conduction

• As mentioned earlier, charge is the key thing
  that must move for electrical currents. In
  physics charge can be an electron, and in
  chemistry we will see electron flow, as well as
  ion flow.
• Both of these occur in typical cells (batteries)

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21-2 Electrodes

• Electrodes are surfaces where reduction or
  oxidation reactions occur
• In a battery, a redox reaction occurs, but
  typically the reduction/cathode is separated from
  the oxidation/anode by a wire and a salt solution
  called a salt bridge
• The separation allows for electron and ion flow
  between the two ½ reactions giving us current

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21-2

• The cathode has the reduction occur, so it gains
  electrons and likewise the anode (oxidation)
  loses electrons. Therefore it only makes sense
  that electrons will flow from the anode to the
  cathode.
• The directionality of electron flow is typically
  1 point on an electrochemistry question on the AP
  exams.

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21-3 The Electrolysis of Molten Sodium Chloride
(The Downs Cell)

• This is a good example to show the basics of
  redox reactions as well as provide a good visual
  example
• Here we have molten NaCl and by running
  electrical current through it, we can produce
  sodium metal and chlorine gas
• Na e- ? Na (l)

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21-3

• Since the sodium ion picks up an electrons or
  Gains an Electron it is Reduced (GER)
• 2Cl- ? 2e- Cl2(g)
• Since the chloride Loses Electrons it is Oxidized
  (LEO)
• Always remember that LEO the lion goes GER

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21-3

• To combine the two ½ reactions, you need to have
  the same number of electrons so that they cancel,
  thus we multiply the Na ½ reaction by 2
• 2x Na e- ? Na (l)
• 2Cl- ? 2e- Cl2(g)
• 2Na 2Cl- ? 2Na(l) Cl2(g)
• Since Na is reduced, it is formed at the cathode
  and Chlorine is formed at the anode

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21-4 The Electrolysis of Aqueous Sodium Chloride

• This reaction is similar, but occurs in solution
  rather than molten sodium chloride
• The redox reaction is the same, but after Na(s)
  is produced, it immediately begins to react with
  water to make Na and H2 and OH-

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21-5 The Electrolysis of Aqueous Sodium Sulfate

• This is actually just an electrolysis of water to
  produce hydrogen and oxygen gas
• While it appears to be a simple reaction, it
  might be a good example to analyze by splitting
  up the net reaction into ½ reactions
• H2O(l) ? H2(g) O2(g)
• The hydrogen in water is a 1 oxidation state, it
  goes to 0 so it is reduced and oxygen is
  therefore oxidized

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21-5

• 2H2O ? O2 4H 4e-
• We can easily now see the loss of electrons for
  this oxidation ½ reaction
• 2H2O 2e- ? H2 2OH-
• Additionally we can see the hydrogen is reduced
  and gains electrons
• To combine these reactions, we would multiply the
  2nd reaction by 2, and the OH- and H would
  combine to form H2O and cancel with 4 waters from
  the reactant side

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21-6 Counting Electrons Coulometry and Faradays
Law of Electrolysis

• In physics, current is measured in Amps, or a
  Couloumb/second, a Couloumb (C) is a measure of
  charge
• In chemistry we can use Amps to determine how
  much electrolysis will occur, similar to
  stoichiometry, using ½ reactions with the of
  electrons as mol ratios

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21-6

• A mol of electrons is called a Faraday and has
  96, 485 C of charge.
• Therefore we can convert between time and amount
  (mass/mol/volume/particles) or material by using
  time, amps, Faradays constant, of electrons
  and the ratios between of electrons and amount
  of substance

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21-7 Commercial Applications of Electrolytic Cells

• Plating of metals is very common, but also
  expensive so typically it is only done as a
  surface and sometimes very thin surfaces. The
  chromium surface of bumpers is only 0.0002 mm
  thick.
• How many atoms thick is that?

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21-8 The Construction of Simple Voltaic Cells

• Voltaic (galvanic) cells have spontaneous redox
  reactions occuring, separated into two halves.
  The half-cells are separated and connected by a
  wire and a salt bridge to maintain neutrality of
  charge and allow charge to flow and be extracted.
  
• A standard cell has all solutions as 1M and all
  gases at 1atm

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21-9 The Zinc-Copper Cell

• In this cell we have a zinc electrode in contact
  with 1M Zinc Sulfate solution and the other ½ has
  a copper electrode in 1M copper sulfate solution
  connected to the zinc by a wire and a salt bridge
• The reaction is
• Cu2(aq) Zn(s) ? Cu(s) Zn2 (aq)

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21-9

• If a statement like the electrode loses mass, or
  gains mass is said, this tells you whether
  reduction or oxidation is occurring at that
  electrode. If metal is plating the electrode to
  make it gain mass, then it is more than likely
  reduction and the cathode
• In the zinc copper cell the copper electrode will
  gain mass and the current flows from the zinc
  electrode to the copper electrode via the wire

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21-9

• The equation for this cell can be simplified to
  ZnZn2 (1M) Cu2 (1M)Cu
• The Double line in the middle indicates a salt
  bridge, the single lines separate the electrodes
  and the ions being reduced or oxidized on those
  electrodes as well as the concentrations.
• The anode is written on the left

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21-9

• If you drop a piece of metal zinc into copper
  sulfate solution, the exact same reaction will
  occur. The point of the salt bridge and the
  separate ½ reactions is to cause a flow of
  electrons from one electrode to another. It is
  the separation that allows us to capture the
  electrical potential energy in these chemicals.

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21-10 The Copper-Silver Cell

• This is another specific example of a cell, we
  wont get into details on this until after we go
  through some thermodynamics and how to find cell
  potentials (voltages)
• There is a reference here back to activity series
  of metals for single displacement reactions way
  back in chapter 4, this is how that list is
  developed.

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21-11 The Standard Hydrogen Electrode

• We cannot determine the potential of a single
  electrode, both reduction and oxidation occur at
  the same time so you cannot just measure one, you
  can only measure one relative to another
• For this reason, we set the standard hydrogen
  electrode (SHE) to be 0 Volts
• From there we can compare electrodes to this one
  and have a set of values to use

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21-11

• The SHE converts protons (H) into hydrogen gas
  (H2) and vice versa. Both reactions are assigned
  a potential of 0 V

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21-12 The Zinc-SHE Cell

• This means that we can compare a zinc electrode
  (ZnZn2 (1M)) in conjunction with the SHE
• This produces a cell with a potential (voltage)
  of 0.763 V and we can therefore assign this
  voltage to zinc
• The voltage of a cell is determined by adding the
  voltage of the anode and cathode

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21-12

• Since the SHE is 0V and the total was 0.763V, the
  zinc electrode must also be 0.763V.
• Zn2 (aq) 2e- ? Zn(s) Eored -0.763V
• Zn(s) ? Zn2 (aq) 2e- Eoox 0.763V
• A spontaneous cell will have a Ecell and so in
  this case the zinc metal will oxidize and the
  hydrogen will be reduced to produce hydrogen gas

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21-13 The Copper-SHE Cell

• Likewise with copper, we can arrange a cell and
  measure the potential difference
• Cu2(aq) 2e- ? Cu(s) Eored .337 V
• Cu(s) ? Cu2(aq) 2e- Eoox -.337 V
• The reverse will always produce the same
  potential with the opposite sign

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21-14 Standard Electrode Potentials

• Now we can use the SHE against all kinds of
  electrode and end up with a list of potentials of
  all electrodes, in addition we can now compare 2
  electrodes where neither is a SHE
• If the electrode has a Eored then it will
  undergo reduction with the SHE and if it is
  then it will undergo oxidation with the SHE

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21-14

• The more positive the value of Eored, the more
  likely that substance is to undergo reduction at
  the cathode in a galvanic cell or the better an
  oxidizing agent it will be
• The reverse is true with large negative values
  for Eored

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21-15 Uses of Standard Electrode Potentials

• The uses of these potentials includes determining
  which direction in a cell will occur
  spontaneously as well as determining the
  potential of that given cell.
• The more positive Eored will be the reaction at
  the cathode being reduced in the spontaneous
  reaction, with zinc and copper we will see copper
  reduced

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21-15

• Cu2(aq) Zn(s) ? Cu(s) Zn2 (aq)
• Cu2(aq) 2e- ? Cu(s) Eored .337 V
• Zn(s) ? Zn2 (aq) 2e- Eoox 0.763V
• Then we can also calculate the Eocell
• Eocell Eored Eoox
• Eocell .337V (0.763V) 1.100 V
• So we will get 1.100 V and electrons will flow
  from the Zn electrode to the Cu electrode

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21-16 Standard Electrode Potentials for Other
Half-Reactions

• A summary of everything you should know by now
• Labeling oxidation, reduction, oxidizing agents,
  reducing agents, cathode, anode, direction of
  flow of electrons
• Calculating the potential of a cell given the net
  reaction
• Going from ½ reactions to a full balanced
  reaction and vice versa

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21-16

• Making calculations involving Faradays Constant
  96485 C and amount of substance deposited during
  electrolysis
• Predicting which metal will be reduced in a
  galvanic cell based on potentials
• Determine spontaneity based on overall cell
  potential
• Describe metals as good oxidizing or reducing
  agents based on reduction potentials

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21-16

• Standard conditions (1M, 1atm, 25 C)
• Describe how cell potentials can be determined
  using a SHE and also know that potentials are
  relative and arbitrarily assigned
• Be familiar with common examples such as
  decomposition of water, Zn/Cu cells and so forth.

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21-17 Corrosion

• Corrosion occurs when metals react with oxygen,
  typically with moisture present
• Corrosion is a big problem because a metal
  surface becomes ionic where it can crumble and
  not have the protection that metals offer
• For iron the reaction is 4Fe(s) 3O2 (g) xH2O
  (l) ? Fe2O3xH2O (s)

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21-18 Corrosion Protection

• To protect from corrosion there are several
  things you can do, interestingly you can add a
  layer of metal on the surface of the iron, and
  the metal can either be a more reactive or less
  reactive metal.
• The less reactive metal will do nothing, but if
  it breaks, the iron will corrode very quickly
  because of the subsequent reaction available

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21-18

• Zinc is more reactive than iron, so galvanized
  steel is often used. The zinc will not rust, and
  if the iron is exposed, the iron will not react
  with oxygen until there is no longer any zinc
  metal in contact.

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21-19 The Nernst Equation

• The Nernst Equation is used for cells that do not
  have standard amounts, for example if I create a
  battery where the concentrations are much higher
  than 1M
• E Eo 0.0592/n logQ
• E is the new potential, n is the number of
  electrons transferred, Q is the reaction quotient
  (from equilibrium), Eo is the standard reduction
  potential

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21-20 Using Electrochemical Cells to Determine
Concentrations

• The Nernst Equation can be used to calculate
  potentials using concentrations, or if the
  potentials are measured, we can measure
  concentration.
• Common examples of this include pH meters which
  use voltage to determine pH in a solution.
• You do not need to know about saturated calomel
  electrodes not glass electrodes for pH meters

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21-20

• The only thing you need to be able to do with the
  Nernst Equation is the calculations of
  concentration and potential

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21-21 The Relationship of Eocell to ?Go and K

• Eocell RTlnK/(nF)
• lnK nFEocell/(RT)
• R is 8.314 J/(molK), n is the of electrons
  transferred, F is Faradays constant (9.65x104 J)
• If you know the cell potential you can calculate
  the equilibrium constant and vice versa

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21-21

• Its important to be able to calculate the cell
  potential and equilibrium constant, but its also
  important to be able to analyze relationships
  between Eocell, ?Go, and K
• For example, if ?Go is -, which means the
  reaction is spontaneous in the forward direction,
  what type of ?Eo would also provide a spontaneous
  forward reaction?

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21-22 Dry Cells 21-25 The Hydrogen-Oxygen Fuel
Cell

• These sections go into specific types of
  batteries and show how some can be recharged and
  others cannot. If youre curious about what
  chemicals are in car batteries or Duracells read
  on. It wont be on the AP test and we wont have
  the time to cover it.