Electrochemistry

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Electrochemistry

• Electrochemistry is a branch of chemistry that
  deals with electrically related applications of
  redox reactions.
• Reduction-oxidation reactions involve the
  transfer of electrons.
• Oxidation means the losing of electrons and
  reduction means the gaining of electrons. The 2
  occur together, they are opposite sides of the
  same coin.
• For example when zinc is in contact with a
  copper II sulfate solution, the zinc strip loses
  electrons to the copper ions in solution. The
  copper ions accept the electrons and fall out of
  solution. As electrons are transferred between
  zinc atoms and copper ions energy is released as
  heat when the reactions are separated we can
  set them up so that instead of heat energy we can
  get electrical energy.

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Example 1

• Net ionic equation
• Zn (s) Cu2 (aq) ? Cu (s) Zn2
  (aq)
• Oxidation
• Reduction

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Voltaic (Galvanic) Cells

• An electrochemical cell, such as a voltaic cell,
  consists of 2 electrodes. Each electrode is in
  contact with an electrolyte. The 2 electrodes are
  connected by a conducting wire or a circuit. And
  a porous barrier separates the 2 half reactions
  (or half cells).
• A voltaic cell specifically deals with a
  spontaneous redox reaction as the source of
  energy. It converts chemical energy into
  electrical energy.

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How it works

• In the wet voltaic cell represented in the
  previous slide an electric current can run
  through an external connecting wire so that the
  electric current moves in closed loop path
  (closed circuit).
• The electrode where oxidation occurs is called
  the anode.
• The electrode where reduction occurs is called
  the cathode.
• The 2 half-reactions occur at the same time but
  in different places at the cell (the porous
  barrier separates them). A salt bridge is
  necessary to keep the half cells electrically
  balanced so that a charge does not build up in
  the cell and stop the electrochemical reaction
  prematurely this salt bridge allows for the
  passage of ions in the cell.

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• ANode, OXidation REDuction, CAThode
• AN OX and a RED CAT

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Zinc Copper
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Cell Notation

• When these cells are represented they are written
  as follows
• Anode electrode anode solution cathode
  solution cathode electrode
• Example 2 Write the cell notation for the
  following reaction
• Zn (s) Cu2 (aq) ? Cu (s) Zn2
  (aq)
• Zn(s)Zn2 (aq) Cu2 (aq) Cu(s)

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Inactive Electrodes
graphite I-(aq) I2(s) H(aq), MnO4-(aq),
Mn2(aq) graphite
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Practice Problems

• Write the half reaction in which I- (aq) changes
  to I2 (s). Identify if this occurs at the anode
  our cathode.
• Nickel solid is oxidized in to Ni2 ions in a
  voltaic cell while the Cu2 ion are being reduced
  in to copper solid atoms.
• Write the half reactions
• Write the net ionic equation
• Identify the anode and cathode
• Write the cell notation

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Cell Voltage

• The cell voltage from a redox reaction is
  referred to as the standard voltage, Eo (unit
  volts, V).
• Standard conditions are 1 atm and 1 M solutions
• Example 3 Zinc metal is placed in hydrochloric
  acid. Zinc is the anode and hydrogen gas forms at
  the cathode. The reaction gives off a standard
  voltage of 0.762 V. Write the net ionic equation
  for this reaction

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Standard Voltage

• When calculating the standard voltage the
  standard voltage from the reduction and oxidation
  reactions must be considered. So when calculating
  the standard voltage you must use the following
  formula
• Eo Eored Eoox

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Standard Potentials

• The standard half cell (half redox reaction)
  voltages are referred to as standard potentials
  and are used to calculate the standard voltage.
• Standard Reduction Potentials
• (one of your equation sheets)
• This table gives the standard reduction
  potentials, the standard oxidation potentials are
  the same magnitude but the reverse sign.

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When Calculating Cell Voltage

• When calculating cell voltages there are 2 main
  points to remember
• 1. the calculation of E , is always a positive
  quantity for a voltaic cell (spontaneous
  reaction).
• 2. The standard cell voltage is independent of
  how the equation for the cell reaction is
  written. This means you must never multiply the
  voltage by the coefficients used to balance the
  chemical equation.

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Example

• Example 4 Use standard reduction potentials to
  calculate the standard voltage for the Zn-H
  cell from example 3.

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• The Pain of a Dental Voltaic Cell
• Have you ever felt a jolt of pain when biting
  down with a filled tooth on a scrap of foil left
  on a piece of food? Heres the reason. The
  aluminum foil acts as an active anode (E of Al
  - 1.66 V), saliva as the electrolyte, and the
  filling (usually a silver/tin/mercury alloy) as
  an inactive cathode. O2 is reduced to water, and
  the short circuit between the foil in contact
  with the filling creates a current that is sensed
  by the nerve of the tooth.

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Finding Eored or Eoox

• If the standard voltage and the cell voltage from
  either the reduction or oxidation reaction the
  other maybe found by a simple rearrangement.
• Example 4 If the standard voltage gathered from
  the standard Zn-Cu2 cell is 1.101 V and the Eoox
  0.762 V, then find the Eored.

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Reducing Oxidizing Agents

• If a species undergoes reduction (gains
  electrons) then it is the oxidizing agent. If it
  undergoes oxidation (loses electrons) then it is
  the reducing agent.
• The stronger the attraction for electrons the
  stronger the oxidizing agent. Or if using the
  standard reduction potentials, the more positive
  the Eored the stronger the oxidizing agent
  (oxidizing strength would be the opposite if
  using the reduction potential table).

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Cell Voltage Gibbs Free Energy Equilibrium

• Standard cell voltage and standard free energy
  are related by the following equation
• Go -nFEo
• When Go lt 0 and Eo gt 0 the reaction is
  spontaneous.
• Standard cell voltage and equilibrium are related
  by the following equation
• Eo RTlnK
• nF
• or at standard conditions (25o C)
• Eo 0.0257 V lnK
• n
• When K gt 1 the reaction is spontaneous

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Effect of Concentration

• Voltage will increase for a reaction if the
  concentration of the reactants is increased or
  that of the products is decreased. This makes the
  reaction more spontaneous.
• Voltage will then decrease if the concentration
  of the reactants is decreased or that of the
  products is increased. This makes the reaction
  less spontaneous.

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The Nernst Equation

• Offers a quantitative relationship between cell
  voltage concentration
• E Eo - RT lnQ
• nF
• or at standard conditions (25oC)
• E Eo - 0.0257 V lnQ
• n

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Interpreting Q

• If Q gt 1 - concentration of the products are high
  so E lt Eo
• (meaning lnQ is positive)
• If Q lt 1 - concentration of the reactants are
  high so E gt Eo
• (meaning lnQ is negative)
• If Q 1 - reaction at standard conditions for
  cell voltage so E Eo
• (meaning lnQ 0)

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Electrolytic Cells

• An electrolytic cell is a non-spontaneous redox
  reaction that made to occur by pumping electrical
  energy into the system.
• When carried out in an electrochemical cell this
  is referred to as electrolysis. This is the
  procedure used when electroplating. Electrons are
  pushed into the cathode and removing them from
  the anode.

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Quantitative Relationships

• Quantitative relationships between the amount of
  electricity passed through an electrochemical
  cell
• For the reaction Cu2 (aq) 2e- ? Cu (s)
• 2 mol e- 1 mol Cu (s) 63.55 g Cu
• Coloumb (C) the quantity of electrical charge
  (or electrical current).
• 1 mol e- 9.648 x 104 C
• Determining Current flow I q/t
• Ampere (A) unit for the rate of current flow,
  (1 A 1 C/s)
• Current (C) over time (s)
• Joule (J) the amount of electrical energy, 1 J
  1 CV

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Electrochemistry Stoichiometry
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From left to right, Walther Nernst, Albert
Einstein, Max Planck, Robert Millikan, and Max
von Laue.
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FRQ 1

• It is observed that when silver metal is placed
  in aqueous thallium(I) fluoride, TlF, no reaction
  occurs. When the switch is closed in the cell,
  the voltage reading is 1.14 V.
• (a) Write the reduction half-reaction that occurs
  in the cell.
• (b) Write the equation for the overall reaction
  that occurs in the cell.
• (c) Identify the anode in the cell. Justify your
  answer.
• (d) On the diagram above, use an arrow to clearly
  indicate the direction of electron flow as the
  cell operates.
• (e) Calculate the value of the standard reduction
  potential for the Tl/Tl half-reaction.
• The standard reduction potential, E, of the
  reaction Pt2 2 e- ? Pt is 1.20 V.
• (f) Assume that electrodes of pure Pt, Ag, and Ni
  are available as well as 1.00 M solutions of
  their salts.
• Three different electrochemical cells can be
  constructed using these materials. Identify the
  two metals that when used to make an
  electrochemical cell would produce the cell with
  the largest voltage. Explain how you arrived at
  your answer.
• (g) Predict whether Pt metal will react when it
  is placed in 1.00 M AgNO3(aq). Justify your
  answer.

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FRQ 2

• 2 H2(g) O2(g) ? 2 H2O(l)
• In a hydrogen-oxygen fuel cell, energy is
  produced by the overall reaction represented
  above.
• (a) When the fuel cell operates at 25C and 1.00
  atm for 78.0 minutes, 0.0746 mol of O2(g) is
  consumed. Calculate the volume of H2(g) consumed
  during the same time period. Express your answer
  in liters measured at 25C and 1.00 atm.
• (b) Given that the fuel cell reaction takes place
  in an acidic medium,
• (i) write the two half reactions that occur as
  the cell operates,
• (ii) identify the half reaction that takes place
  at the cathode, and
• (iii) determine the value of the standard
  potential, E, of the cell.
• (c) Calculate the charge, in coulombs, that
  passes through the cell during the 78.0 minutes
  of operation as described in part (a).

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FRQ 3
An external direct-current power supply is
connected to two platinum electrodes immersed in
a beaker containing 1.0 M CuSO4(aq) at 25C, as
shown in the diagram above. As the cell operates,
copper metal is deposited onto one electrode and
O2(g) is produced at the other electrode. The two
reduction half-reactions for the overall reaction
that occurs in the cell are shown in the table
below.
Half-Reaction E0(V)
O2(g) 4 H(aq) 4 e- ? 2 H2O(l) 1.23
Cu2(aq) 2 e- ? Cu(s) 0.34
(a) On the diagram, indicate the direction of
electron flow in the wire. (b) Write a balanced
net ionic equation for the electrolysis reaction
that occurs in the cell. (c) Predict the
algebraic sign of ?G for the reaction. Justify
your prediction. (d) Calculate the value of ?G
for the reaction. An electric current of
1.50 amps passes through the cell for 40.0
minutes. (e) Calculate the mass, in grams, of the
Cu(s) that is deposited on the electrode. (f)
Calculate the dry volume, in liters measured at
25C and 1.16 atm, of the O2(g) that is produced.