Electrochemistry: Chemical Change and Electrical Work

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Electrochemistry Chemical Change and Electrical
Work
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REVIEW

• Oxidation the loss of electrons by a species,
• leading to an increase in oxidation number of one
  
• or more atoms
• Reduction the gain of electrons by a species,
• leading to an decrease in oxidation number of one
  
• or more atoms
• Oxidizing agents the species that is reduced in
  a
• redox reaction

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• Reducing agents the species that is oxidized in
  a redox reaction

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In acidic solution add H or H2O only In basic
solution add OH- or H2O only
 
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THE HALF-REACTION METHOD

• This method breaks the overall reaction into its
• two components half-reactions
• Each half-reaction is balanced separately and
• then added
• Use the following guidelines to help
• Write as much of the unbalanced net ionic
• equation as possible
• 2. Decide which atoms are oxidized and which are
• reduce write the two unbalanced half-reactions

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• 3. Balance by inspection all atoms in each half-
• reaction except H and O
• 4. Use the rules mentioned previously to balance
• H and O in each half-reaction
• 5. Make equal the number of electrons involved in
• both half-reactions
• Take a look at the breathalyzer reaction
• H(aq) Cr2O72-(aq) C2H5OH(l) ?
• Cr3(aq) C2H4O(l) H2O(l)

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Balance the following net ionic equation in
basic solution.
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ELECTROCHEMISTRY

• Deals with chemical changes produced by an
• electric current and with the production of
• electricity by chemical reactions
• All electrochemical reactions involve transfer
• of electrons and are redox reactions
• EChem reactions take place in electrochemical
• cell (an apparatus that allows a reaction to
• occur through an external conductor)

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ELECTROCHEMICAL CELLS
Two types 1. Electrolytic cells - these are
cells in which an external electrical source
forces a nonspontaneous reaction to occur 2.
Voltaic cells - also called galvanic cells. In
these cells spontaneous chemical reactions
generate electrical energy and supply it to an
external circuit
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• Electric current enters and exits the cell by
• electrodes - electrodes are surfaces upon which
• oxidation or reduction half-reactions occur
• Inert electrodes - electrodes that dont react
• Two kinds of electrodes
• 1. Cathode - electrode at which reduction
• occurs (electrons are gained by a species)
• 2. Anode - electrode at which oxidation occurs
• (as electrons are lost by some species)

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VOLTAIC CELLS

• Cells in which spontaneous reactions produces
• electrical energy
• The two half-cells are separated so that
  electron
• transfer occurs through an external circuit
• Each half-cell contains the oxidized and reduced
• forms of a species in contact with each other
• Half-cells linked by a piece of wire and a salt
• bridge

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• A salt bridge has three functions
• 1. It allows electrical contact between the two
  
• half-cells
• 2. It prevents mixing of the electrode
  solutions
• 3. It maintains electrical neutrality in each
• half-cell as ions flow into and out of the salt
  
• bridge
• Point 2 is important no current would flow if
• if both solutions were in the same cell

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• Point 3 is also important anions flow into the
  
• oxidation half-cell to counter the build-up
• of positive charge
• Current flow spontaneously from negative to
• the positive electrode
• In all voltaic cells the anode is negative and
  the
• cathode is positive

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• In voltaic cells, voltage drops as the reaction
• proceeds. When voltage 0, the reaction is at
• equilibrium

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The Silver-Copper cell

• Composed of two half-cells
• 1. A strip of copper immersed in 1 M CuSO4
• 2. A strip of silver immersed in 1 M AgNO3
• Experimentally we see
• - Initial voltage is 0.46 volts
• - The mass of the copper electrode decreases
• - The mass of the silver electrode increases
• - Cu2 increases and Ag decreases

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• Cu ? Cu2 2e- (oxidation, anode)
• 2(Ag e- ? Ag) (reduction, cathode)
• 2Ag Cu ? Cu2 Ag (Overall cell reaction)
• Cu Cu2(1.0 M) Ag(1.0 M) Ag
• Notice that in this case the copper electrode is
  
• the anode

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STANDARD ELECTRODE POTENTIALS

• Associated with each voltaic cell is a potential
• difference called the cell potential, Ecell
• E measures the spontaneity of the cells redox
• reaction
• Higher (more positive) cell potentials indicate
  a
• greater driving force for the reaction as written
• All electrode potentials are measured versus the
• Standard Hydrogen Electrode (SHE) E 0.00 V

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• The Ecell calculated is for the cell operating
• under standard state conditions
• For electrochemical cell standard conditions
• are
• solutes at 1 M concentrations
• gases at 1 atm partial pressure
• solids and liquids in pure form
• All at some specified temperature, usually 298 K

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• The electrode potential for each half-reaction
  is
• written as a reduction process
• The more positive the E value for a half-
• reaction the greater the tendency for the
  reaction
• to proceed as written
• The more negative the E value, the more likely
• is the reverse of the reaction as written

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Prediction of Spontaneity
1. First write the HR equation with the more
positive (less negative) E for the reduction
along with its potential 2. Write the other HR
as an oxidation and include its oxidation
potential 3. Balance the electron transfer 4. Add
the reduction and oxidation HR and add the
corresponding electrode potentials to get
the overall cell potential, Ecell
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• Important points to note
• 1. E for oxidation half-reactions are equal to
• but opposite in sign to reduction half-reactions
• 2. Half-reaction potentials are the same
• regardless of the species stoichiometric
• coefficient in the balanced equation
• Ecell gt 0 Forward reaction is spontaneous
• Ecell lt 0 Backward reaction is spontaneous

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Ecell, ?G and K

• From thermodynamics, we know that,
• DG -RT lnK
• We can relate Ecell to free energy for that
  cell
• DG -nFEcell
• n number of moles of e-
• So -nFEcell -RT lnK and
• Ecell (RT/nF) lnK

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• (Standard state conditions)
• Under nonstandard conditions
• ?G -nFEcell

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THE NERNST EQUATION

• Usually concentrations of reactants differ from
• one another and also change during the course
• of a reaction
• As cell reaction proceeds, cell voltage drops so
  
• that Ecell is different from Ecell
• Ecell and Ecell are related by the Nernst
• Equation
• Ecell Ecell - (RT/nF) lnQ

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Ecell Ecell - (RT/nF) lnQ E potential
under the nonstandard conditions E standard
potential R gas constant, 8.314 J/mol.K T
absolute temperature n number of moles of
electrons transferred F faraday, 96,485 J/V.mol
e- Q reaction quotient
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BATTERIES

• Two type of batteries
• - Primary batteries cannot be recharged
• Once all the chemicals are consumed there is
• no more chemical reaction
• - Secondary batteries can be regenerated
• Most common example is the lead storage
• battery used to power automobiles

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The Lead Storage Battery

• Composed of two alternating groups of Pb
• plates one group contains pure lead (anode) and
• the other group contains PbO2 (cathode)
• The plates are immersed in 40 sulfuric acid
• During discharge
• Pb ? Pb2 2e- (oxidation)
• Pb2 SO42- ? PbSO4 (precipitation)
• Net Pb SO42- ? PbSO4 2e- (anode)

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• At the cathode
• PbO2 4H 2e- ? Pb2 2H2O (reduction)
• Pb2 SO42- ? PbSO4 (precipitation)
• Net reaction
• PbO2 4H SO42- 2e- ? PbSO4 2H2O
• Adding the HR for the two half-cells, gives
• Pb PbO2 4H 2SO42- ? 2PbSO4 2H2O
• Ecell 2.041 V
• The battery can be recharged

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• Fuel Cells
• These are galvanic cells in which the reactants
• are continuously supplied to the cell and the
• products are continuously removed
• Best known example is the hydrogen-oxygen
• fuel cell
• Hydrogen is fed into the anode compartment
• and oxygen into the cathode compartment

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• Oxygen is reduced at the cathode porous
• carbon doped with metallic catalysts
• At the anode hydrogen is oxidized to water
• Anode 2H2(g) 4OH-(aq) ? 4H2O(l) 4e-
• Cathode O2(g) 2H2O(l) 4e- ? 4OH-(aq)
• Overall 2H2(g) O2(g) ? 2H2O(g)

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CORROSION

• Ordinary corrosion is a redox process in
• which metals are oxidized by oxygen in the
• presence of moisture
• A point of strain on the surface of the metal
• acts as an anode
• Areas on the metal surface exposed to air
• serves as cathodes

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• Anode Fe(s) ? Fe2(aq) 2e-
• Cathode O2(g) 4H(aq) 4e- ? 2H2O(l)
• 4Fe(s) O2(g) 4H(aq) ?
• 4Fe2(aq) 2H2O(l)
• 2Fe2(aq) 4H2O(l) ? Fe2O3H2O(s) 6H
• Rust
• Al also undergo corrosion initial oxidation
• is stopped by a layer of Al2O3

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Corrosion prevention 1. Plating a metal with a
thin layer of a less easily oxidized metal 2.
Allow a protective film to form naturally on the
surface of the metal 3. Galvanizing coating
the metal with zinc 4. Cathodic protection
connecting the metal to a sacrificial anode
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ELECTROLYTIC CELLS

• Cells in which an electric current causes a
• nonspontaneous reaction to occur one common
• process is called electrolysis
• In electrolytic cells the anode is the positive
• electrode and the cathode is the negative
• electrode
• Still Anode oxidation cathode reduction

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• The Down Cell Electrolysis of molten NaCl
• Using graphite inert electrodes the following
• observations are made
• 1. Chlorine, Cl2, is liberated at one electrode
• 2. Sodium metal forms at the other electrode
• Explanation
• 1. Chlorine is produced at the anode by the
• oxidation of Cl- ions

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2. Metallic sodium is formed by reducing Na
ions at the cathode
2Cl- ? Cl2(g) 2e- (oxidation, anode
HR) 2(Na e- ? Na(l) (reduction, cathode
HR) 2Na 2Cl- ? 2Na(l) Cl2(g) Overall cell
rxn.
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• Electrons used at the cathode are reproduced at
• the anode
• The reaction is nonspontaneous and electricity
• is used to force the reaction to occur

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• Electrolysis of aqueous sodium chloride
• In an EChem cell containing aqueous NaCl
• - H2 gas is liberated at one electrode
• - Cl2 gas is liberated at the other electrode
• - Solution at the cathode is basic
• Rationalization
• - Chloride ions are oxidized at the anode and
  
• H2O is reduced at the cathode

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• 2Cl- ? Cl2 2e- (oxidation, anode)
• 2H2O 2e- ? 2OH- H2 (reduction, cathode)
• 2H2O 2Cl- ? 2OH- H2 Cl2 Overall
• Sodium metal is more active than hydrogen
• metal and liberates H2 from solution
• The hydroxide ions are responsible for the
• basicity around the cathode

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FARADAYS LAW

• States that the amount of substance that
• undergoes oxidation or reduction at each
• electrode during electrolysis is directly
• proportional to the amount of electricity that
• passes through the cell
• One faraday the amount of electricity that
• reduces or oxidizes 1 equivalent of a substance

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• One equivalent of any substance is the amount
• of that substance that supplies or consumes one
• mole of electrons
• 1F 1mole of electrons
• 6.022 x 1023 e-
• 96,485 C
• C It
• C charge passed I current
• t time (in seconds)