Ch 17a Electrochemistry

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Ch 17a Electrochemistry

• Redox Review
• Electrochemical cells
• Electrochemistry, equilibrium and thermodynamics
• Electrolysis
• Corrosion prevention
• pH Meters

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OxidationReduction Reactions

• Identify which reactant is
• Being oxidized? Reduced?
• the reducing agent and which is the oxidizing
  agent?
• Balance the redox equation
• Sn2(aq) Fe3(aq) ? Sn4(aq) Fe2(aq)
• Ca(s) H(aq) ? Ca2(aq) H2(g)
• SnO2(s) C(s) ? Sn(s) CO(g)

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Balancing Redox Reactions

• Half-Reaction Method focus on the transfer of
  electrons.
• The reaction is broken into half-reactions and
  then recombined
• important when considering batteries and
  electrochemistry.
• Determine oxidation and reduction half-reactions
• Balance each ½-Rxn for atoms
• Balance each ½-Rxn for charge by adding e
• Combine and cancel to form one equation
• Review Balancing Redox in acidic or basic soln

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Redox Reactions and Spontaneity

• Zn(s) Cu2(aq) ? Zn2(aq) Cu(s)
• BUT
• Zn2(aq) Cu(s) ? No Reaction
• Oxidation Half-Reaction Zn(s) ? Zn2(aq) 2
  e
• Reduction Half-Reaction Cu2(aq) 2 e ? Cu(s)

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Electrochemical Cell Potentials

• The larger positive reduction half reaction from
  Table 18.1 will be the reduction rxn.
• The standard potential (total) for any galvanic
  cell is
• Ecell Eoxidation Ereduction
• Standard half-cell potentials are always quoted
  as a reduction rxn.
• The sign must be REVERSED for the oxidation rxn.
• Referenced against the standard hydrogen
  electrode (S.H.E.)
• a platinum electrode in contact with H2 gas (1
  atm) and aqueous H ions (1 M).
• assigned an arbitrary value of exactly 0.00 V

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Practice

• Determine the standard cell potential (EMF) for
• Sn2(aq) Fe3(aq) ? Sn4(aq) Fe2(aq)
• The reverse reaction?
• Which is spontaneous?

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Electrochemical Cells (galvanic cells)

• Electrodes
• metal strips connected by a wire.
• Anode oxidation takes place.
• Cathode reduction takes place.
• electrons flow from anode to cathode
• NOTE Current, I, is opposite direction (cathode
  to anode) in physics/electronics
• Salt Bridge tube that contains a solution of an
  inert electrolyte.

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Electrochemical Cell Notation

• Conventions for writing the cell
• Electrode Anode Soln Cathode Soln
  Electrode
• Pt(s) H2 (1 atm) H (1 M) Fe3(aq),
  Fe2(aq) Pt(s)

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Electrochemical Cells
Pt wire
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Electrochemical Cells

• Identify the oxidation and reduction ½-rxn anode
  and cathode
• Write the cell notation
• Calculate the EMF
• of an electrochemical cell made of a Cd electrode
  in a 1.0 M Cd(NO3)2 solution and a Cr electrode
  in a 1.0 M Cr(NO3)2 solution?
• of an electrochemical cell made of a Mg electrode
  in a 1.0 M Mg(NO3)2 solution and a Ag electrode
  in a 1.0 M AgNO3 solution?

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Electrochemistry, equilibrium and thermodynamics
combined (just run !!!!)

• ?G RT ln K
• ?G nFEcell
• nFEcell RT ln K
• n moles of e- transferred
• F 96,500 coulombs

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Just run faster

• nFEcell RT ln K
• is rearranged at 25C to give
• Ecell (0.0257?n) ln K
• or
• Ecell (0.0592?n) log K

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Spontaneity of a Reaction

• Calculate the standard free energy change (?G)
  and the equilibrium constant (K) for the
  following reactions at 25C
• Sn(s) 2 Cu2(aq) ?? Sn2(aq) 2 Cu(aq)
• Fe2(aq) 2 Ag(s) ?? Fe(s) 2 Ag(aq)
• 4 Fe2(aq) O2(g) 4 H(aq) ?? 4 Fe3(aq) 2
  H2O(l)

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Variations from standard conditions-the Nernst
Equation

• Variations from 25 C and 1.0 M solutions (1 atm
  gases) change the cell potentials to non-standard
  values
• ?G ?G RT lnQ
• nFE nFE RT lnQ
• Nernst Equation

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The Nernst Equation

• Metallic copper with iron(III) to give copper(II)
  and iron(II). What is the potential of a cell
  when Fe3 0.0001 M, Cu2 0.25 M, and
  Fe2 0.20 M?
• Metallic zinc with hydrochloric acid. Calculate
  the cell potential at 25C when H 1.0 M,
  Zn2 0.0010 M, and PH2 0.10 atm.

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Electrolysis

• Electrolysis Uses Current to drive a
  non-spontaneous reaction to completion
• purification of metals or production of gases
• electroplating

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Electrolysis of Water

• H2O (l) ? H2 (g) O2 (g)
• Anode Water is oxidized to oxygen gas.
• 2 H2O(l) ? O2(g) 4 H(aq) 4 e
• Cathode Water is reduced to hydrogen gas.
• 4 H2O(l) 4 e ? 2 H2(g) 4 OH(aq)

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Quantitative Electrolysis

• Amount of chemical electrolyzed depends on the
  quantity of electrons (charge) passed through the
  cell.
• To produce 1 mole of Na (l) requires how many
  moles of e-?
• To produce 1 mole of Cl2 (g) requires?

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Electrolysis

• To determine the moles of electrons
• Charge (C) Current (A) x Time (s)
• Because the charge on 1 mol of e is 96,500 C
• A constant current of 30.0 A is passed through an
  aqueous solution of NaCl for a time of 1.00 hour.
  How many grams of Na and how many liters of Cl2
  gas at STP are produced?

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Corrosion

• Corrosion is the oxidative deterioration of
  metal.
• 25 of steel produced in USA goes to replace
  steel structures and products destroyed by
  corrosion.
• Rusting of iron requires the presence of BOTH
  oxygen and water.
• Rusting results from tiny galvanic cells formed
  by water droplets.

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Corrosion

• Oxidation Fe(s) ? Fe2(aq) 2 e
• Reduction O2(g) 4 H(aq) 4 e ? 2 H2O(l)
• Overall
• 2 Fe(s) O2(g) 4 H(aq) ? 2 Fe2(aq) 2 H2O(l)

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Corrosion Prevention

• Galvanizing
• Coat the iron (steel) with zinc.
• Zinc is more easily oxidized than iron,
• which protects the iron how?
• And reverses the oxidation of the iron.

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Cathodic Protection

• Cathodic Protection
• is the protection of (iron) from corrosion by
  connecting it to a metal (a sacrificial anode)
  that is more easily oxidized.
• sacrificial anode - usually magnesium or zinc.

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pH and the Nernst Equation

• A important use of the Nernst equation is the pH
  electrode.
• Consider a hydrogen electrode (anode) and a
  reference cathode placed in a solution of unknown
  pH
• Pt H2 (1 atm) H (? M) Reference Cathode
• Ecell EH2 ? H Eref
• The Nernst equation can be applied to the
  half-reaction
• H2(g) ? 2 H(aq) 2 e

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pH and the Nernst Equation

• E 0 V PH2 is 1 atm n 2

A higher cell potential indicates a higher pH,
therefore we can measure pH by measuring Ecell.
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pH and the Nernst Equation

• Ag(s) AgCl(s) HCl(aq) glass H(aq)
  reference
• The difference in H from one side of the glass
  membrane to the other causes a potential to
  develop, which adds to the measured Ecell.
• Overall cell potential is
• A higher cell potential indicates a higher pH,
  therefore we can measure pH by measuring Ecell.

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pH and the Nernst Equation

• A glass electrode (Ag/AgCl wire in dilute HCl)
  with a calomel reference is the most common
  arrangement.
• Ag(s) Cl(aq) ? AgCl(s) e E 0.22 V
• Calomel Hg2Cl2(s) 2 e ? 2 Hg(l) 2 Cl
  E 0.28 V

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pH and the Nernst Equation

• The following cell has a potential of 0.55 V at
  25C
• Pt(s) H2 (1 atm) H (? M) Cl (1 M)
  Hg2Cl2(s) Hg(l)
• What is the pH of the solution at the anode?
• The following cell has a potential of 0.28 V at
  25C
• Pt(s) H2 (1 atm) H (? M) Pb2 (1 M)
  Pb(s)
• What is the pH of the solution at the anode?