Electrochemistry

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Chapter 21

• Electrochemistry

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Oxidation-Reduction Reactions

• oxidation loss of electrons LEO
• reduction gain of electrons GER
• Can be determined from change in oxidation
  numbers.

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Oxidation Number

• Charge an atom has or would have if all the
  bonding electrons were assigned to the most
  electronegative element.

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Rules for assigning oxidation numbers

• 1). Oxidation number pure neutral element 0
  (e.g. Li, H2, C etc)
• 2). Oxidation number monatomic ion ionic charge
  (Na 1, Mg2 2, etc)
• 3). Oxidation number of O -2 in all its
  compounds except
• those with O-O bond (peroxides e.g. Na-O-O-Na
  then ox. number -1) superoxides, (e.g. KO2 then
  ox. number -1/2)

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Rules cont.

• 4). H is 1 in all its compounds (except metal
  hydrides then -1)
• 5). F is - 1 in all its compounds (except, of
  course, F2)
• 6). Alkali metals are always 1 alkaline earths
  are always 2.
• 7). Sum of oxidation number is zero for neutral
  compounds, and equal to the overall charge for
  polyatomic ionic species.

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• Oxidation and Reduction must occur simultaneously
• Zn(s) Cu2(aq) -------gt Zn2(aq) Cu(s)

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Example 1

• Indicate which of the reactants is reducing and
  which is oxidizing
• 2Ce4(aq) Sn2(aq) ? 2Ce3(aq) Sn4(aq)
• 8H(aq) MnO4-1(aq) 5Fe2(aq) ? 5Fe3(aq)
  Mn2(aq) 4H2O(l)

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Electrochemical Cells

• System consisting of electrodes that dip into an
  electrolyte and in which an oxidation-reduction
  reaction either uses or generates an electric
  current.

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Types of Electrochemical Cell

• Voltaic or Galvanic
• A spontaneous reaction generates an electric
  current
• Electrolytic
• An electric current is used to drive a
  nonspontaneous reaction

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Electrodes

• Cathode electrode where reduction reaction
  occurs
• Anode electrode where oxidation reaction
  occurs.

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Balancing Redox Reactions

• Neutral Solution
• 1. Split into two half-reactions
• 2. Balance each half-reaction
• a. Balance the elements
• b. Balance the charge
• 3. Combine half-reactions so that electrons
  cancel.

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Example 2

• Cl2(g) Br-(aq) --gt Br2(l) Cl-(aq)

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Balancing Redox Reactions

• Acidic Solution
• 1. Split into two half-reactions
• 2. Balance each half-reaction
• a. Balance the elements
• b. Balance oxygens by adding water
• c. Balance hydrogens by adding hydrogen ions
• d. Balance the charge
• 3. Combine half-reactions so that electrons
  cancel.

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Example 3

• Zn(s) VO2(aq) --gt V2(aq) Zn2(aq)

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Balancing Redox Reactions

• Basic Solution
• 1. Split into two half-reactions
• 2. Balance each half-reaction
• a. Balance the elements
• b. Balance oxygens by adding water
• c. Balance hydrogens by adding hydrogen ions
• d. Neutralize the H ions by adding OH- ions
• e. Balance the charge
• 3. Combine half-reactions so that electrons
  cancel.

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Example 4

• Zn(s) Cl2 (g) --gt Zn(OH)2(s) Cl- (aq)

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Voltaic Cell

• Description
• anode Zn ----------gt Zn2 2e-
• cathode Cu2 2e- ----------gt Cu
• Consists of two half-cells that are electrically
  connected. The salt bridge consists of a tube of
  an electrolyte that is connected to two
  half-cells of the voltaic cell. Allows the flow
  of charge ions but prevents diffusional mixing of
  the different solutions.

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Voltaic Cell Key Points

• 1. There are always two separate half-cells
  connected by a wire and a salt bridge
• 2. One of the compartments is the anode and the
  other is the cathode
• 3. If a metal participates in a cell reaction it
  is ordinarily chosen as the electrode. If no
  metal is involved in the half-reaction an
  electrically conduction solid like platinum or
  graphite is used
• 4. Electrons flow from the anode to the cathode.
• 5. Cations move to the cation and anions to the
  anode via the salt bridge.

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Cell Notation

• Zn/Zn2 ll Cu2/Cu
• anode reaction is shown at left
• salt bridge is indicated by ll
• cathode reaction is shown at right

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Example 5

• Draw cell diagram and write the cell notation
• Fe3(aq) H2(g) ? Fe2(aq) 2H(aq)

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Electrochemical Cells and Potentials

• Emf
• Electrons generated at the site of oxidation of a
  cell are thought to be "driven" or "pushed"
  toward the cathode by an electromotive force, or
  emf. This force is due to the difference in
  electric potential of an electron at the two
  electrodes.

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Standard Potential

• Measure of the driving force of the cell reaction
  when all ions and molecules in the solution are
  at a concentration of 1 M and all gases are at a
  pressure of one atm.

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Calculating the Potential Eo of an
Electrochemical Cell

• Eotot Eoox Eored
• Eoox standard voltage for the oxidation
  half-reaction
• Eored standard voltage for the reduction
  half-reaction

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Standard Hydrogen Half-Cell

• Cannot determine an individual half-reaction
  voltage therefore we arbitrarily take the
  standard voltage for the reduction of H ions to
  H2(g) to be zero (standard hydrogen half-cell)
• 2H(aq,1M) 2e- ----------gt H2(g, 1atm) Eored
  0
• With this assigned voltage others can be
  determined from measurement by hook-up with
  knowns.

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Standard Hydrogen Half-Cells
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Standard Potentials
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Using Standard Potentials

• 1. The Eo values are for reactions written in
  the form "oxidized form electrons --------gt
  reduced form." The species on the left side of
  the reaction is an oxidizing agent, and the
  species on the right is a reducing agent. All
  potentials are therefore for reduction reactions.
• 2. When writing the reaction "reduced form
  --------gt oxidized form electrons," the sign
  of Eo is reversed, but the value of Eo is
  unaffected.
• 3. All the half-reactions are reversible.
• 4. The more positive the value of Eo for the
  reactions the better the oxidizing ability of the
  ion or compound.

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• 5. The more negative the value of the reduction
  potential Eo the less likely the reaction occurs
  as a reduction, and the more likely the reverse
  reaction occurs.
• 6. The reaction between any substance on the
  left in this table with any substance lower than
  it on the right is product-favored under standard
  conditions.
• 7. The algebraic sign of the half-reaction
  potential is the sign of the electron when it is
  attached to H2/H3O standard cell.
• 8. Electrochemical potential depends on the
  nature of the reactants and products and their
  concentrations, not on the quantities of material
  used.

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• Large negative value means oxidation strongly
  favored strong reducing agent.
• Large positive value means reduction strongly
  favored strong oxidizing agent.
• Relative values in table give an indication that
  one half-reaction favored over other. Summing
  half-cell reactions allow determination of
  standard cell potential.
• Half-cell potential intensive property ?
  independent of amount of material ? we dont use
  stoichiometric coefficients for determining
  standard cell potentials.

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Example 6

• Determine the cell potential of
• Br2(l) 2I?(aq) ? I2(l) 2Br?(aq)

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Example 7

• Determine the cell potential of
• 2Ag(aq) Cu(s) ? 2Ag(s) Cu2(aq)

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Example 8

• Determine cell potential
• MnO4-(aq) Fe(s) ? Fe2(aq) Mn2(aq)
  (balanced?)
• when it is operated galvanically. Which is the
  oxidizing agent? reducing agent?

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Example 9

• Determine if the reaction below is spontaneous in
  the direction written.
• Fe3(aq) Ag(s) ? ?

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Eotot and DGo

• 1. DG free energy change
• maximum amount of useful work
• -wmax
• 2. Electric work
• charge x potential energy difference
• coulomb x volt

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Eotot and DGo

• wmax nFE coulombs x volts electrical energy
  in joules where F is the Faraday constant, 9.65 x
  104 J/V mol.
• DGo -nFEo tot -96.5 nEo tot (in KJ)
• Because spontaneous reactions have a negative
  free energy change, DG, the negative sign in the
  equation above confirms that all product-favored
  electron transfer reactions have a positive Eo.

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Electrochemical Cells at Nonstandard Conditions

• Voltage is a measure of reaction spontaneity.
• Hence
• voltage is increased by increasing concentrations
  of reactants or decreasing concentrations of
  products.
• voltage is decreased by decreasing concentrations
  of reactants or increasing concentrations of
  products.

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Quantitative Nernst Equation

• aA bB ----------gt cC dD
• Etot Eotot -RT/nF ln products/reactants
  raised to the power of their coefficients
• where R 8.314510 J/K mol, F 9.645309 x 104
  J/v mol and number of electrons transferred at 25
  C

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Nernst Equation

• Etot Eotot - 0.0257/n lnproducts/reactants
  raised to the power of their coefficients
• n number of electrons transferred
• use molarities for aqueous solutions, atm for
  gases

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K Equilibrium Constant

• Etot Eotot - - 0.0257/n ln products/reactants
  raised to the power of their coefficients
• at equilibrium Etot zero
• and
• K products/reactants raised to the power of
  their coefficients
• therefore
• 0 Eotot - 0.0257/n ln K

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Eotot and K

• ln K nEotot/0.0257

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Example 10

• Determine free energy and equilibrium constant
  for reaction below (unbalanced).
• MnO4-(aq) Fe(s) ? Fe2(aq) Mn2(aq)

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Example 11

• Determine cell potential and equilibrium constant
  of Cl2/Br2 cell.

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Example 12

• The following cell has a potential of 0.578 V at
  25C determine Ksp.
• Ag(s)AgCl(s)Cl?(1.0 M)Ag(1.0 M)Ag(s).

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Batteries

• A battery is a package of one or more galvanic
  cells used for the production and storage of
  electric energy by chemical means. A galvanic
  cell consists of at least two half cells, a
  reduction cell and an oxidation cell. Chemical
  reactions in the two half cells provide the
  energy for the galvanic cell operations.
• Each half cell consists of an electrode and an
  electrolyte solution. Usually the solution
  contains ions derived from the electrode by
  oxidation or reduction reaction.

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Types of Batteries

• Batteries can be divided into two types primary
  or disposable batteries and secondary or
  rechargeable batteries.

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• Fuel cells are different from batteries in that
  they consume reactant, which must be replenished,
  whereas batteries store electrical energy
  chemically in a closed system.

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Corrosion of Metals/Cathodic Protection

• Cathodic protection (CP) is a technique to
  control the corrosion of a metal surface by
  making that surface the cathode of an
  electrochemical cell.

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Cathodic Protection
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Electrolytic Cell
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Aqueous Electrolysis

• Cathode(reduction)
• a. cation reduced to metal usually occurs with
  transition metal cations
• b. H ions reduced to H2 occurs with strong
  acids
• c. H2O molecules reduced to H2 occurs with
  Group 1, Group 2 metals and Al
• 2H2O(l) 2e- ----------gt H2(g) 2OH-(aq)

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Aqueous Electrolysis

• Anode(oxidation)
• a. anion oxidized to nonmetal
• b. OH- ions oxidized to O2 occurs with strong
  bases
• c. H2O molecules oxidized to O2 with NO3-,
  SO42- and F-

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Overall result

• NiCl2 Ni Cl2
• NaCl H2, OH-, Cl2
• CuSO4 Cu, O2, H

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Example 13

• What reactions take place at the anode and
  cathode when each of the following is
  electrolyzed?
• 1.0 M NiBr2
• 1.0 M AlF3
• 1.0 M MnI2

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Stoichiometry of Electrolysis

• Units
• Faraday charge 9.65 x 104 C 1 mole of e-
• Coulomb charge
• Ampere current charge/time C/s
• Joule electrical work C x volts
• watt energy/time 1 joule/second
• volt potential difference

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• Calculations
• Electric Charge electric current x time
  elapsed
• Electric Energy Coulombs x volts
• watt volts x amperes

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Example 14

• Determine amount of Cu2 electrolyzed from
  solution at constant current of 6.00 A for period
  of 1.00 hour.

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Example 15

• What volumes of H2(g) and O2(g) rat STP are
  produced from the electrolysis of water by a
  current of 2.50 A in 15.0 minutes?

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Example 16

• Electrolysis of a molten metal chloride (MCl3)
  using a current of 6.50 amp for 1397 seconds
  deposits 0.471 g of metal at the cathode. What
  is the identity of the alkaline earth metal?