Cells

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Cells

• Electrochemical Cell Contains electrodes
  dipping into electrolyte solutions. The
  half-reactions occur at two different electrodes
  and a flow of electrons (current) is involved.
• Voltaic (galvanic ) cells a spontaneous redox
  reaction occurs and causes a flow of electrons.
• Electrolytic cells an external source of
  electric current is used to drive an otherwise
  nonspontaneous redox reaction.

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Voltaic (Galvanic) Cells

• Consider the spontaneous redox reaction that
  occurs if a piece of copper is immersed in a
  solution of silver nitrate (occurring in one
  beaker)
• Cu(s) 2 Ag(aq) ? Cu2(aq) 2 Ag(s)
• No useful work is done in this process (e.g., no
  usable electric current). Some heat will be
  released to the surroundings.
• How can we harness this spontaneous reaction in
  such a way as to yield useful work?
• Answer Separate the half-reactions in two
  containers.

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Figure 17.2Galvanic Cells
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• Ex Suppose we have a piece of Zn metal dipping
  into a 1 mol/L solution of ZnSO4 in the left-hand
  compartment and a piece of Cu metal dipping into
  a 1 mol/L solution of CuSO4 in the right-hand
  compartment.
• If we connect the metal electrodes with a wire,
  the following observations are made
• The piece of Zn decreases in mass over time.
• The piece of Cu gains mass over time.
• The concentration of the Zn2(aq) in the
  left-hand compartment increases over time.
• The concentration of Cu2(aq) in the right-hand
  compartment decreases over time.

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• What is the salt bridge?
• It contains a solution of an electrolyte such as
  NaNO3 in a gel, such that the ions are mobile but
  will not mix rapidly with the other solutions.
• If we place a voltmeter between the two
  electrodes, we obtain a reading of 1.10 V.
• Now, lets analyze more closely what is happening
  in the cell.

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Figure 17.6 A Zn/Cu Galvanic Cell
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• At the left-hand electrode (anode half-reaction)
• Zn(s) ? Zn2(aq) 2 e oxidation
• At the right-hand electrode (cathode
  half-reaction)
• Cu2(aq) 2 e ? Cu(s) reduction
• Add them to get the overall reaction
• Zn(s) Cu2(aq) ? Zn2(aq) Cu(s)
• But now the electrons lost by the Zn anode flow
  through the external circuit to the Cu cathode
  (an electric current which we can make use of)

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• Note By convention the anode is placed on the
  left and the cathode on the right when picturing
  a galvanic cell.
• You should be able to describe the half-reactions
  at each electrode and the direction of movement
  of all species (including electrons and the ions
  of the salt bridge).

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Inert Electrodes

• Sometimes a half-reaction will not involve a
  solid substance
• MnO4-(aq) 8H (aq) e ? Mn2 (aq) 4 H2O
  (l)
• In such cases, an inert metal such as Pt is used
  as an electrode to which the wire is attached. A
  example is shown in the following figure.

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Voltaic Cells with Inert Electrodes

• In this example the one half reactions has no
  solid metal so carbon solid is used as the
  electrode.
• Do the analysis of this cell.

C(s)
Zn(s)
Porous Cup
Electrolyte
MnO4(aq) H (aq)
Zn 2(aq)
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Figure 17.7A Schematic of a Galvanic Cell
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• Voltaic Cells have two electrodes

• Cathode electrode
• (GERC)
• Gain of
• Electrons is the process of
• Reduction and occurs at the
• Cathode.

• Anode electrode
• (LEOA)
• Loss of
• Electrons is the process of
• Oxidation and occurs at the
• Anode.

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Cathode vs. Anode

• Cathode
• GERC
• Is a solid.
• Is the location of reduction.
• Electrons are on the reactant side.
• Electrons are consumed.
• Has a charge.

• Anode
• LEOA
• Is a solid
• Is the location of oxidation.
• Electrons are on the product side.
• Electrons are produced.
• Has a charge.

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• Cell Potential (or EMF or voltage)
• Think of this as the push that sends electrons
  from the anode to the cathode.
• We measure (with a voltmeter) the potential
  difference between the anode and cathode.
• ?Ecell Ecathode Eanode
• The superscript refers to standard-state
  conditions and reversible cell operation.

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• ?Ecell must be positive in order for a
  spontaneous reaction to occur. Thus, the cathode
  must be at a higher potential than the anode
  (think of electrons rolling downhill).
• Ex for the Zn-Cu cell previously described, ?E
  1.10 V.

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• The Standard Hydrogen Electrode (S.H.E.)
• 2 H(1 mol/L) 2e ? H2(1 atm) E 0.00
  V
• is used as the reference for all half-reactions

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Analysis of a Voltaic Cell

• Write the two half cells including the voltages
  (Change the sign on the oxidation half-reaction
  from the one listed in your data booklet).
• Balance electrons and write the net reaction with
  the net voltage. (dont change the voltage)
• Identify the cathode. (GERC)
• Identify the anode. (LEOA)

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Analysis of a Voltaic Cell (contd)

• Net voltage is the sum of the reduction potential
  and oxidation potential.
• Identify the electrodes and electrolytes.
• Identify the directions of electron movement in
  the wire.
• Identify the movement of the anions and cations
  in the solution.

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Voltaic Cell Cu(s) / Cu2(aq) // Zn 2(aq) /
Zn(s)
Electrode
Salt Bridge
Zn(s)
Cu(s)
Cu2(aq)
Zn 2(aq)
Electrolyte
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A Zn/H Galvanic Cell
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2 H(aq) 2e- H2(g)
Zn (s) Zn2(aq) 2e-
2 H(aq) Zn (s)
H2 Zn2(aq)
The voltaic cell on the previous slide is fully
described with the following notation
Zn(s) Zn(aq) H(aq), H2(g) Pt(s) )
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Electrons flow from the anode to the cathode in a
voltaic cell. (They flow from the electrode at
which they are given off to the electrode at
which they are consumed.) Reading from left to
right, this line notation therefore corresponds
to the direction in which electrons flow.
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Line Notation For Voltaic Cells

• Voltaic cells can be described by a line notation
  based on the following conventions.
• Single vertical line indicates change in state or
  phase.
• Within a half-cell, the reactants are listed
  before the products.
• A double vertical line indicates a junction
  between half-cells.
• The line notation for the anode (oxidation) is
  written before the line notation for the cathode
  (reduction).

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• In a galvanic cell, the electrode with the higher
  E value will be the cathode.
• Ex In a cell using the Ni2Ni (0.25 V) and
  AgAg (0.80 V) half-reactions, Ag will be the
  cathode
• and the spontaneous reaction will be
• Ni(s) 2 Ag(aq) ? Ni2(aq) 2 Ag(s)
• ?E 0.80 (0.25) 1.05 V
• Important Notice that the sign of the potential
  of the anode is used as found in the table (not
  reversed). This is because we subtracted the
  anode potential from the cathode potential.

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Voltaic Cells

• Primary Cells
• Produce energy until one component is used up,
  then discarded
• Examples are the zinc chloride cell and the 9
  volt
• Secondary Cells
• Store energy and may be recharged
• Examples are Ni-Cad cells and lead acid car
  batteries

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Dry Cells

• The ordinary zinc carbon cell
• Anode (oxidation)
• Zn (s) ? Zn 2 (aq) 2e
• Cathode (reduction)
• 2MnO2 (s) NH4 (aq) 2e ? Mn2O2 (s) 2NH3
  (aq) H2O (l)

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Dry Cells

• A new cell produces about 1.5V
• Once reaction reaches equilibrium its flat

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Dry Cell
Metal Cap ()
Mixture of Carbon Manganese Dioxide
CathodeCarbon Rod
Ammonium Chloride Zinc Chloride Electrolyte
Anode Zinc Case ()
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Alkaline Cells (eg. Duracell)

• The ordinary zinc carbon cell
• Anode (oxidation)
• Zn (s) ? Zn 2 (aq) 2e
• Immediately reacts with OH ions in the
  electrolyte to form zinc hydroxide
• Zn (s) 2OH (aq) ? Zn(OH)2 (s) 2e

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Alkaline Cells

• Cathode (reduction)
• 2MnO2 (s) H2O(l) 2e ? MnO2 (s) OH (aq)
  H2O (l)
• Five times the life of the dry cell

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Alkaline Cell
Metal Cap ()
Cathode outer steel case
Potassium Hydroxide Electrolyte
Powdered Zinc
AnodeSteel or Brass
Mixture of Carbon Manganese Dioxide
Metal Base ()
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Button Cells

• Used in very small applications like watches,
  cameras etc.
• Two main types
• Mercury zinc and silver zinc
• Anode (Oxidation)
• Zn (s) 2OH (aq) ? Zn(OH)2 (s) 2e

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Button Cells

• Cathode Reduction ()
• depends on the type of battery
• HgO(s) H2O (l) 2e ? Hg (l) 2OH (aq)
• Ag2O(s)H2O (l) 2e ? 2Ag (s) 2OH (aq)
• Produce an almost constant 1.35V

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Button Cell
Metal Cap ()
Zinc Powder
Cathode outer container of nickel or steel ()
Electrolyte
Mercury Oxide
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Secondary Cells

• Lead Acid (Car Battery)
• Nickel cadmium Cells
• Fuel Cells

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Lead Acid Battery

• Car Batteries
• Also called storage batteries or accumulators
• Each cell produces 2 volts so typical 12 volt car
  battery contains 6 cells
• Both electrodes are lead plates separated by some
  porous material like cardboard

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Lead Acid Battery

• Positive electrode is coated with PbO2 Lead (IV)
  Oxide
• The electrolyte is a solution of 4M sulfuric acid

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Lead Acid Battery

• Anode Oxidation ()
• Pb(s) SO4 2- ? PbSO4 (s) 2e
• Cathode Reduction ()
• PbO2(s) SO4 2- 4H 2e ? PbSO4 (s) 2H2O
  (l)
• Overall Reaction
• Pb(s) PbO2(s) 2H2SO4 ? 2PbSO4 (s) 2H2O (l)

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Nickel Cadmium Cells

• Often called Nicads
• Electrodes are nickel and cadmium
• Electrolyte is potassium hydroxide
• Reactions involve the hydroxides of the two metals

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Nickel Cadmium Cells

• Anode (Oxidation)
• Cd (s) 2OH (aq) ? Cd(OH)2 (s) 2 e
• Cathode (Reduction)
• NiO-OH (s) H2O (l) e ? Ni(OH)2 (s) OH
  (aq)
• Overall Reaction
• Cd (s) NiO-OH(s) H2O(l) ? Cd(OH)2 (s)
  Ni(OH)2 (s)

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Fuel Cells

• Limitation of dry cells looked at so far is that
  they contain reactants in small amounts and when
  they reach equilibrium.
• Primary Cells are then discarded, secondary cells
  are then recharged
• A cell that can be continually fed reactants
  would overcome this and allow for a continual
  supply of electricity

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Fuel Cells

• Fuel cells transform chemical energy directly
  into electrical energy
• 60 efficiency
• Space Program uses hydrogen and oxygen with an
  electrolyte of potassium hydroxide

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Fuel Cells

• Anode (Oxidation)
• H2(g) 2OH (aq) ? 2H2O (l) 2e
• Cathode (Reduction)
• O2(g) 2H2O(l) 4e ? 4OH(aq)
• Overall Equation
• H2(g) O2(g) ? 2H2O (l)

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Hydrogen Oxygen Fuel Cell


Electrolyte
Oxygen Gas Inlet
HydrogenGas Inlet
Porous Anode
Porous Cathode
Water outlet
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Some basic generalizations about Voltaic Cells

• The net voltage is always positive.
• The cathode is the location of reduction. (GERC)
• The anode is the location of oxidation.(LEOA)
• Electrons always move from the anode to the
  cathode.
• Cations move to the cathode, anions move to the
  anode.

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Electrolytic Cells

• In an electrolytic cell, a non-spontaneous redox
  reaction is forced to occur by applying an
  external voltage to the cell
• The potential for an electrolytic cell is
  negative. This is the voltage that must be
  supplied to cause the reaction to occur.

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• This is a diagram of a Hoffman apparatus which is
  used to separate water into hydrogen and oxygen
  by applying an electrical current.

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• Predict the products in the electrolysis of water
• What evidence would confirm your prediction?

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Electrolysis of Water

• Prediction
• Cathode
• 2H2O 2e- H2 2OH- E-0.83V
• Anode
• 2H2O O2 4H 4e- E-1.23V
• Net
• 2H2O 2H2 O2 E -2.06V

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Electron flow in the circuit is opposite to the
conventional current flow. The reaction at the
cathode (tube A) is the reduction of water2
H2O(aq) 2e- H2 (g) 2 OH-(aq)Oxidation
takes place at the anode (tube B). The oxidation
of water according to the reaction 2 H2O (l)
4 H(aq) O2( g) 4e-has a standard
electrode potential of -1.23V, so the overall
reaction is therefore2 H2O (l) 2 H2 (g) O2
(g)
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• Predict the products for the electrolyis of brine
  (aqueous sodium chloride)

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Strange Salt

• In an aqueous solution of sodium chloride, water
  would be predicted to be both the SOA and the SRA
  so the predicted reaction in an electrolytic cell
  is the same as that found for the electrolysis of
  water
• 2 H20 (l) 2 H2 (g) O2 (g)

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Strange Salt (contd)

• However, the aqueous chloride ion is oxidized
  preferentially before the water, so the two half
  reactions and net reaction are
• 2 Cl-(aq) Cl2 (g) 2e-
• 2 H2O(aq) 2e- H2 (g) 2 OH-(aq)
• 2 Cl-(aq) 2 H2O(aq) Cl2 (g) 2 OH-(aq)

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Applications of Electrolysis

• Production of reactive metals from their molten
  salts (eg. Al from Al2O3 Na from NaCl)
• Electroplating objects to enhance their value
  and/or corrosion resistance
• Electrorefining of metals

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Aluminum production
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Electroplating
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Electrorefining of Copper
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Voltaic vs. Electrolytic Cells

• Spontaneous
• Converts chemical energy to electrical energy
• Oxidation occurs at the anode reduction occurs
  at the cathode

• Non-spontaneous
• Converts electrical energy to chemical energy
• Oxidation occurs at the anode reduction occcurs
  at the cathode

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Voltaic vs. Electrolytic Cells

• Positive potential
• SOA and SRA react
• SOA is above SRA on table
• Electrons move from anode to cathode
• Two separated half cells are used

• Negative potential
• SOA reacts with the SRA
• Electrons forced to move from anode to cathode by
  external power source
• Generally, only one container is used

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Cell Stoichiometry

• The amount (moles) reduced or oxidized in a cell
  (whether voltaic or electrolytic) is directly
  proportional to the moles of electrons
  transferred.

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Cell Stoichiometry (contd)

• The mathematical relationship is

where I current (amperes
or Coulombs/second) t time (seconds) F
Faradays constant 96500 Coulombs/molee
ne It
F
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Cell Stoichiometry (contd)

• What is the mass of copper that can be plated out
  in an electrolytic cell if 4.7 Amperes are
  delivered for 2.1 hours?
• Known I 4.7 C/s
• t 2.1 h x 3600 s/h 7560 s
• F 96500 C/mole
• Unknown moles of electrons (which will lead to
  moles of copper

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Cell Stoichiometry

• ne (4.7 C / s) x (7650 s)
• 96500 C / mole
• 0.37 mole
• molCu 0.37 mole x 1 molCu x 63.55 g
• 2 mole 1 mol
• 12 g Cu

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Corrosion and Electrochemistry

• Corrosion is simply the oxidation of structures
  by environmental substances
• In a relatively neutral environment, the OA is
  the H2O / O2 combination
• O2(g) H2O(l) 4e- 4 OH-(aq)

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Corrosion and Electrochemistry

• In an acidic environment, the OA is the H /O2
  combination
• O2(g) 4 H(aq) 4e- 2 H2O(l)
• Notice that this OA is higher on the redox table,
  so it is a more powerful OA

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Protection From Corrosion

• The cost of corrosion means that steps must be
  taken to minimize or eliminate its effects
• One of the easiest ways is to coat the metal
  object (ie. Painting or coating with oil)
• Another way is to galvanize the metal (this is a
  coating of zinc metal on the object

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Protection From Corrosion

• Cathodic protection (also called sacrificial
  anode) is a means of protecting an expensive
  metal object by sacrificing an inexpensive metal
• Consider ships (made of steel, which is mainly
  iron), floating on salt water, a wonderful
  electrolyte. Why dont they corrode and sink to
  the bottom of the ocean?

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Protection From Corrosion (contd)

• The reason they do not is because of cathodic
  protection (sacrificial anodes)
• A plate of zinc or magnesium metal is attached on
  the inside hull of the ship. When the OA (O2/H2O)
  takes electrons from the iron atoms at the
  surface of the ship, electrons are conducted from
  the zinc to replace the stolen electrons. In this
  way, the iron is preserved in its elemental state
  and the zinc is sacrificed (notice on the redox
  table that zinc is a stronger RA than iron so it
  will react instead)