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Ch.17 Outline
A. Electrochemistry - the study of the interchange between chemical and electrical energy.
a. Examples from life
- rusting
- calculators/computers
- batteries/solar panels
1. Galvanic Cells
a. redox reactions - involve the transfer of electrons from the reducing agent to the oxidizing agent
- oxidation - involves the loss of electrons (increase in oxidation state)
- reduction - involves the gaining of electrons ( decrease in oxidation state)
b. half-reactions - show either reduction or oxidation :
- e.g. Reduction : H+ + MnO4- + 5e- Æ Mn2+ + 4 H2O
Oxidation : 5(Fe2+ Æ Fe3+ + e-)
c. galvanic cell - uses a spontaneous redox reaction which converts chemical energy into electrical
energy
- salt bridge (or porous disk) - allows the exchange of ions to prevent the build up of charges
in each half-cell
- anode - electrode in the cell where oxidation occurs
- cathode - electrode in the cell where reduction occurs
- Cathodes and anodes are collectively called electrodes, if a solid metal is used in a galvanic
cell it can be used as an electrode. If an aqueous or gaseous substance is oxidized or
reduced, an inert electrode must be use (usually a platinum or carbon rod).
b. cell potential (Ecell) (or electromotive force - emf) - the push on the electrons which causes
them to flow
- volt (V) - unit of electrical potential
- one volt equals 1 joule per coulomb of charge transferred (1 V = 1 J/C)
- measured by a voltmeter (lacks accuracy due to resistance within the meter or a
potentiometer - powered by an outside source to avoid inaccuracies of voltmeter)
2. Standard reduction potentials
a. reduction of hydrogen ( 2 H+ + 2e- Æ H2 ) in its standard state ( [PH2 = 1 atm and [H+] =1 M]
is assigned the value of 0.00 V (called standard hydrogen electrode)
b. standard reduction potential of other substances in their standard state are determined by
comparison to the standard hydrogen electrode
e.g. to calculate the standard reduction potential of Zn (Zn2+ + 2e- Æ Zn) a galvanic cell
would be constructed using hydrogen and zinc in their standard states to measure the total
potential of the cell (Ecell) and the standard reduction potential of zinc would be calculated
by the equation :
Eºcell = EºH+Æ H2 + EºZn2+ ÆZn
º symbol indicates standard state (1 atm for gases, 1.0 M for aqueous solutions)
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- if a half-reaction is reversed, (making it an oxidation half-reaction) its sign is reversed
- in a pair of substances, the substance with the largest standard reduction potential will be
reduced (cathode reaction) and the substance with the smaller standard reduction
potential will be oxidized at the anode (reaction will need to be reversed and sign
changed).
- Eºcell is an intensive property so that the magnitude of Eºcell is not changed even when
multiplied by an integer to balance the reaction
c. calculation of cell potential
d. line notation for a galvanic cell :
e.g. oxidation half reaction : Mg Æ Mg2+ + 2ereduction half reaction : Al3+ + 3 e- Æ Al
- the line notation will be : Mg(s) |Mg2+(aq) || Al3+(aq) | Al(s)
- note oxidation is first, reduction is second and the line notation follows the flow of electrons,
electrons flow from Mg to Al as Mg is oxidized to Mg2+ and aluminum is reduced from Al3+
to Al
e. a cell will always run spontaneously in the direction that produces a positive cell potential
3. Cell potential, Electrical Work and Free Energy
a. the potential difference or emf is measured in volts where 1 V = 1 joule per coulomb or :
emf = potential difference (V) = work (J) / charge (C) or E = -w/q
- w is negative because the system is doing the work
- rearranging the above gives -w = qE or
wmax = -qEmax
- remember actual work obtained will always be less than Emax due to the second law of
thermodynamics
b. efficiency of the cell = w/wmax x 100% where w is the work actually obtained from a cell
c. cells and electron flow
- q represents quantity of charge and equals nF when n =number of moles and F = a faraday
which has the value of 96,485 coulombs per mole therefore wmax = -nFEmax
d. cells and free energy
- since ∞G = wmax (from ch. 16), ∞G = -nFE and ∞Gº = -nFEº
4. Dependence of Cell Potential on Concentration
a. Le Chatelier's principle - cell potential will shift as reactants and products shift due to changes
in concentration
b. concentration cells - cells that have the same substances, but at different concentrations
(difference in concentration causes the flow of e-)
c. Nernst equation
E = Eº - ((RT/nF)lnQ)
- the cell will continue to produce electricity until equilibrium is reached at which point Q = K
and Eºcell = 0 and both sides of the cell have the same free energy (∞G = 0)
d. Ion-Selective Electrodes
- electrodes that are sensitive to concentrations of certain ions
e. Calculation of Equilibrium Constants for Redox Reactions
at equilibrium E cell = 0 and Q = K
5. Batteries
a. battery - a galvanic cell or group of galvanic cells connected in series
b. lead storage battery - six cells in a series, each produces about 2 volts giving 12 volts total
- anode is solid lead
- cathode is solid lead coated with lead dioxide
- the electrolyte is sulfuric acid
anode (oxidation) :
Pb + HSO4- Æ PbSO4 + H+ + 2ecathode (red.)
:
PbO2 + HSO4- + 3H+ + 2e- Æ PbSO4 + 2H2O
Cell reaction :
Pb + PbO2 + 2H+ + 2HSO4- Æ PbSO4 + 2 H2O
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From Zumdahl, Chemistry, 4th ed. p. 841
Forward reaction produces the electricity to start the car, run lights, radio etc.. The reverse
direction requires the input of electricity from the alternator which "charges" the battery restores the original reactants.
c. Other batteries
- common dry cells - produce about 1.5 V
- acid version of common dry cell :
- anode reaction :
Zn Æ Zn2+ + 2e- cathode reaction : 2NH4+ + 2MnO2 + 2e- Æ Mn2O3 + 2NH3 + H2O
- basic (alkaline) version - lasts longer due to zinc not being corroded as fast by acidic
conditions
- anode reaction :
Zn + 2OH- Æ ZnO + H2O + 2e- cathode reaction :
2MnO2 + H2O + 2e- Æ Mn2O3 + 2OH-
From Zumdahl, Chemistry, 4th ed. p. 842
From Zumdahl, Chemistry, 4th ed. p. 843
d. Fuel cells - a galvanic cell in which the reactants are continuously supplied
- e.g. the hydrogen-oxygen fuel cell
- anode :
2H2 + 4OH- Æ 4 H2O + 4e- cathode :
4e- + O2 + 2H2O Æ 4 OH- anode and cathode are porous carbon rods containing catalysts
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6. Corrosion
a. corrosion - the oxidation of metals - spontaneous process (positive Eº value) that often returns
metals to their ores
- metals oxidize easily because their standard reduction potentials are less positive than oxygen
- noble metals - gold, silver, platinum, copper - difficult to oxidize and often found in pure state
- silver and copper will tarnish - forming silver sulfide or copper carbonate (greenish
color)
- many metals (e.g. aluminum) form a protective oxide coating which prevents further
oxidation
b. corrosion of iron
- nonuniformities due to composition or stress in steel create regions where oxidation tends to
occur (anodic regions) leaving the surrounding areas cathodic (areas where reduction tends to
occur)
- anodic regions :
Fe Æ Fe2+ + 2e- cathodic regions : O2 + 2H2O + 4e- Æ 4OH- Fe2+ ions must travel from anodic region to cathodic region through moisture (electrolyte)
- in the cathodic region the Fe2+ ion reacts with oxygen to form rust which is a complex of
hydrated iron (III) oxide :
4Fe2+ + O2 + (4 + 2n)H2O Æ 2Fe2O3 ⏐ nH2O + 8 H+
- because of the migration of the ions rust can form a distance from the source of the Fe2+
ions leaving pits and weakness in steel
- salt speeds up the rusting process because of its electrolytic nature and the fact that the
chloride ion forms a stable complex with the Fe3+ ion
c. protection of iron
- galvanizing - the use of zinc to cover steel
- zinc is more active than iron so it tends to oxidize (corrode) and in the process reduce
iron
- alloying - stainless steel contains chromium and nickel both of which tend to form protective oxide
coatings that give stainless steel a reduction potential similar to that of the noble metals
- cathodic protection - attaching a more reactive metal to iron (e.g. magnesium) which is a
better reducing agent and therefore supplies electrons to iron keeping it reduced (in the form
of Fe)
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7. Electrolysis
a. electrolytic cell - a cell used for electrolysis (opposite of galvanic cell)
b. electrolysis - using electricity to force a cell to produce a chemical change for which a cell potential
is negative (nonspontaneous)
- the amount of change depends on the amount of current and the length of time
- unit : ampere = 1 coulomb per second (Cs-1 or C/s)
c. calculations - stoichiometry problems
- current and time ℑ quantity of charge in coulombs ℑ moles e- ℑ moles reactant or product ℑ
mass
d. electrolysis of water - an electrolyte is needed because pure water is a nonelectrolyte
2H2O Æ 2 H2 + O2
e. electrolysis of a mixture of ions
- in a mixture of ions the ions with the most positive Eºcell value will be reduced first and the ions
can be separated according to these values
- e.g. A mixture of Cu2+ (Eºcell = .40 V), Cr3+ (Eºcell = -0.73 V), and Zn2+ (Eºcell = -0.76 V). The
Cu2+ will separate first and then the Cr3+ and finally the Zn2+.
- Overvoltage - sometimes the amount of voltage actually needed is greater than the Eºcell
values predict - called overvoltage. Therefore using the above method to predict order of
separation must be used with caution.