Electrochemistry

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Electrochemistry

• Voltaic cell
• Standard reduction potential
• Electrochemical Series

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Voltaic cell
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• Name Hamza Eid chemistry department
• Title of Lesson Voltaic cell
• Date of Lesson 18 26/3/ - 2007 
• Length of Lesson 4- blocks
• Description of Class Level 11 Advance Chemistry

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Overview

• This lesson allows students to learn how to
• set up voltaic cells and explore the
• electrochemistry that is involved. Students
• will practice writing half reactions, learn to
• designate the cathode and anode, and make
• a reduction table.

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Outcomes

• Students will be able to
• Setup, conduct, and explain voltaic cell
  reactions
• Write equations for half-reactions and balanced
  full redox reactions
• Identify the anode and cathode
• Identify the type of reactions that occur at the
  anode and cathode

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Teacher Does

• Demo - How to set up a half-cell
• The half cell consists of the oxidized and
  reduced form of an element or species.
• In our case a solid metal and a solution of its
  ion. (i.e Zn and Zn2)
• The half cell will be constructed in a 100 mL
  beaker with 50 mL of solution and a strip of the
  metal.
• Two half cells are used to make a voltaic cell.
•  

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• The voltaic cell needs a salt bridge connecting
  the two half cells.
• Our salt bridge will consist of strips of paper
  soaked in a brine (saturated salt) solution.
•   Two half cells can be connected via a wire.
  This experiment will simply use the wire that
  connects the cells to the voltammeter.

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• The voltammeter measurers the potential as
  electrons move through the cell. It will measure
  negative or positive flow depending on which way
  the wires are connected. Be sure to take positive
  measurements only.

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Student Does

• Students will be divided into groups of 3-4
• Ask students, in groups (each group consist of
  3-members) to do the following activity (activity
  1) then each group demonstrates
• How to set up an electrochemical cell using two
  metal /ion half-cells.
• How they can identify the anode and the cathode
  in a cell

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Activity 2- Voltaic Cells Lab

• Resources, materials and supplies needed
• For each group (3-4 students)
• 1 Voltammeter (able to read voltage below 1V)
• 20 mL potassium chloride Solution
• 10 - strips of filter paper
• 6 - 100 mL beakers
• 50 mL of each of the following 1M solutions
  ZnSO4, FeSO4, NiSO4, CuSO4, Ag2SO4, Al2(SO4)3
• 1 - 2" long wire of each of the following wires
  Zn, Fe, Ni, Cu, Ag, Al

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• Supplementary materials, handouts
• Lab Sheets
• Discussion Questions

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• Students will be divided into groups of 3-4 and
  given a worksheet that outlines the procedure.
• For each voltaic cell students will report the
  half reaction equations, the balanced full
  equation, the anode and the cathode, the sites of
  oxidation and reduction, the positive voltage.

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Activity-3You can get electricity from a lemon.

• Make two slits in the skin of a lemon and insert
  a piece of aluminum and a piece of copper in the
  other slit.

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• Make sure the two metals are not touching each
  other inside the lemon. If you touch your tongue
  to the two strips of metal you will feel a tingle
  of electricity.
• The current flows because of the chemical
  reaction that takes place between the metals and
  acid in the lemon juice. The lemon juice acts in
  the same way as Volta's salt water or the
  chemical paste in a battery

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• Assessments

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• Examine the following combination of half -cells.
• Cd(s)Cd2(aq) half cell combined with Ag
  (aq)Ag(s)
• Pt(s) IO3-(aq), H(aq) half cell combined with
  Zn2(aq)Zn(s)
• Pb(s)Pb2(aq) half cell combined with
  Ni2(aq)Ni(s)
• C(s)ClO4-(aq), H(aq) , Cl- (aq) half cell
  combined with Fe3(aq)Fe(s)
• C(s)SO42-(aq), H(aq),H2SO3(aq) half cell
  combined with Pt2(aq)Pt(s)

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For each of the above create a sketch of the cell
similar to the following
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Vocabulary

• Battery - a cell that produces an electrical
  charge from a chemical reaction
• Electrodes - conductors of an electrical charge
• Current - a continuous flow of electrical charges
  
• Voltage - potential difference in a chemical cell
  which produces current
• Anode - negative electrode in a chemical cell
• Cathode - positive electrode in a chemical cell

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• Standard electrode potentials

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Outcomes

• Define standard reduction potential and cell
  potential
• Construct, and label the parts of a standard
  hydrogen electrode
• predict the voltage (E ) of an electrochemical
  cell using the table of standard reduction
  half-cells

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Teacher Does

• Standard Reduction Potentials The tendency of a
  half reaction to occur as reduction under
  standard conditions
• standard conditions
• Temperature 25oC
• concentrations of all ions 1 M
• pressure of all at gases 1 atm)

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• To measure relative electrode potentials, we must
  establish an arbitrary standard.
• That standard is the Standard Hydrogen Electrode
  (SHE).
• The SHE is assigned an arbitrary voltage of
  0.000000 V
• Hydrogen Half-Cell
• reversible reaction

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Hydrogen Electrode

• The standard hydrogen electrode consists of a
  piece of platinum foil coated with fine particles
  of platinum.
• It is immersed into a solution of hydrogen ions
  and hydrogen gas is bubbled over it.
• .

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Cell Potential

• the potential difference, in volts, between the
  electrodes of an electrochemical cell
• positive value indicates a spontaneous reaction

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The Copper-SHE Cell

• The cell components are
• A Cu strip immersed in 1.0 M copper (II) sulfate.
• The other electrode is a Standard Hydrogen
  Electrode.
• A wire and a salt bridge to complete the circuit.
• The initial cell voltage is 0.337 volts.

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The Copper-SHE Cell

• In this cell the SHE is the anode
• The Cu2 ions oxidize H2 to H.
• The Cu is the cathode.
• The Cu2 ions are reduced to Cu metal.

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The Zinc-SHE Cell

• For this cell the components are
• A Zn strip immersed in 1.0 M zinc (II) sulfate.
• A wire and a salt bridge to complete the circuit.
• The other electrode is the Standard Hydrogen
  Electrode.
• The initial cell voltage is 0.763 volts.

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The Zinc-SHE Cell

• The cathode is the Standard Hydrogen Electrode.
• In other words Zn reduces H to H2.
• The anode is Zn metal.
• Zn metal is oxidized to Zn2 ions.

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• In an electrochemical cell. we can think that in
  an electrode the following situation occurring
• The better the reducing agent the more the
• equilibrium lies over the right hand side.
  

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Student Does

• Ask students, in groups (each group consist of
  3-members) to do the following activity (activity
  2)
• Each group demonstrate their observation.
• Use the following work sheet

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• From the activity we conclude that
• If we had used a voltmeter in the circuit we
  would find that the bigger voltage would be
  developed for Zn/Ag.
• the second biggest by Zn /Cu
• and the smallest by Cu/Ag.
• So we can see that the voltage in a cell is
  governed the tendency of the metal ions to go
  into solution represented by the following
  equilibrium.

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• So by measuring the voltage (and the direction of
  flow of the electrons) in a series of cells we
  will be able to arrange oxidising and reducing
  agents in an electrochemical series.

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• The activity series of metals

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Teacher does

• electrochemical series a series of half
  reactions arranged in the ordering of their
  reducing ability.
• The metals are arranged in the order of their
  reducing power.

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• With the most reactive metals at the top. A
  metal placed in solution can displace a metal ion
  of a metal below it in the reactivity series so
• Zn(s) CuSO4(aq) ?ZnSO4(aq) Cu(s)
  
• Will proceed

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Teacher does - Activity-4

• Ask students, in pairs, to design an
  investigation to determine the order of
  reactivity of the metals provided by constructing
  a potential difference chart

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Activity -5

• From Video 1
• Ask how they could improve the design of their
  investigation.

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Activity-6

• Provide students with data of standard electrode
  potentials for a range of half-cells.
• Do a worked example to show how to determine the
  direction of reaction for each half-cell, and how
  to write a balanced equation for the reaction and
  determine the standard cell potential by
  combining the two relevant standard electrode
  potentials.
• Ask students to work in pairs and challenge each
  other to do the same with a different set of
  half-cells and check their partners answer

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• conclusions

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Uses of Standard Electrode Potentials

• Electrodes that force the SHE to act as an anode
  are assigned positive standard reduction
  potentials.
• Electrodes that force the SHE to act as the
  cathode are assigned negative standard reduction
  potentials.
• Standard electrode (reduction) potentials tell us
  the tendencies of half-reactions to occur as
  written

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• For example, the half-reaction for the standard
  potassium electrode is
• The large negative value tells us that this
  reaction will occur only under extreme conditions

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• Compare the potassium half-reaction to fluorines
  half-reaction

• The large positive value denotes that this
  reaction occurs readily as written.
• Positive E0 values denote that the reaction tends
  to occur to the right.
• The larger the value, the greater the tendency to
  occur to the right.
• It is the opposite for negative values of Eo.

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• Uses of Standard Electrode Potentials

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• Use standard electrode potentials to predict
  whether an electrochemical reaction at standard
  state conditions will occur spontaneously.
• Example 21-3 Will silver ions, Ag, oxidize
  metallic zinc to Zn2 ions, or will Zn2 ions
  oxidize metallic Ag to Ag ions?

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• Steps for obtaining the equation for
• the spontaneous reaction.
• Choose the appropriate half-reactions from a
  table of standard reduction potentials.
• Write the equation for the half-reaction with the
  more positive E0 value first, along with its E0
  value.

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• Add the reduction and oxidation half-reactions
  and their potentials. This produces the equation
  for the reaction for which E0cell is positive,
  which indicates that the forward reaction is
  spontaneous

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• Write the equation for the other half-reaction as
  an oxidation with its oxidation potential, i.e.
  reverse the tabulated reduction half-reaction and
  change the sign of the tabulated E0.
• Balance the electron transfer.

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• Example
• Will silver ions, Ag, oxidize metallic zinc to
  Zn2 ions, or will Zn2 ions oxidize metallic Ag
  to Ag ions?

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• Example
• Will permanganate ions, MnO4-, oxidize
• iron (II) ions to iron (III) ions, or will iron
• (III) ions oxidize manganese(II) ions to
• permanganate ions in acidic solution?
• Follow the steps outlined in the previous slides.
• Note that E0 values are not multiplied by any
  stoichiometric relationships in this procedure.

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• Example
• Will permanganate ions, MnO4-, oxidize iron (II)
• ions to iron (III) ions, or will iron (III) ions
• oxidize manganese(II) ions to permanganate ions
• in acidic solution?

• Thus permanganate ions will oxidize iron (II)
  ions to iron (III) and are reduced to manganese
  (II) ions in acidic solution.

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• Example
• Will nitric acid, HNO3, oxidize arsenous acid,
• H3AsO3, in acidic solution? The reduction
• product of HNO3 is NO in this reaction.

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Assessments

• Standard reduction potential
• Test traning

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resourcesSuggested Website Recourse

• http//www.chemguide.co.uk/inorganic/redoxmenu.htm
  ltop20
• http//www.chem.iastate.edu/group/Greenbowe/sectio
  ns/projectfolder/flashfiles/electroChem/voltaicCel
  l20.html
• http//www.chem.iastate.edu/group/Greenbowe/sectio
  ns/projectfolder/animations/SHZnV7.html

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• http//www.chem.iastate.edu/group/Greenbowe/sectio
  ns/projectfolder/animations/SHECu.html
• http//www.wwnorton.com/chemistry/tutorials/ch1
  7.htm
• http//www.chem.iastate.edu/group/Greenbowe/secti
  ons/projectfolder/flashfiles/electroChem/voltaicCe
  llEMF.html

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• http//www.wwnorton.com/chemistry/tutorials/ch17.h
  tm
• http//www.chem.iastate.edu/group/Greenbowe/sectio
  ns/projectfolder/flashfiles/electroChem/electrolys
  is10.html