1. Chapter 19 Electrochemistry
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2. 19.1 Balancing Redox Reactions
Half-reaction method – acid solution
Can add H2O
Can add H+
Steps
Separate the unbalanced reaction into half-reactions.
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Balance each of the half-reactions with regard to atoms other than O and H.
Balance both half-reactions for O by adding H2O.
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Balance both half-reactions for H by adding H+.
Balance both half-reactions for charge by adding electrons.
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multiply one or both of the half-reactions
by the number(s) required to make the number of electrons the same in both.
add the balanced half-reactions together and cancel any identical terms that appear on both sides.
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Additional steps needed in basic solution
For each H+ ion in the final equation, add one OH ion to each side of the equation, combining the H+ and OH ions to produce H2O.
Make any additional cancellations made necessary by the new H2O molecules.
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19.2 Galvanic Cells
Galvanic cell - the experimental apparatus for generating electricity through the use of a spontaneous reaction
Electrodes
Anode (oxidation)
Cathode (reduction)
Half-cell - combination of container, electrode and solution
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Salt bridge - conducting medium through which the cations and anions can move from one half-cell to the other.
Ion migration
Cations – migrate toward the cathode
Anions – migrate toward the anode
Cell potential (Ecell) – difference in electrical potential between the anode and cathode
Concentration dependent
Temperature dependent
Determined by nature of reactants
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cell diagram convention
anode
cathode
salt bridge
phase boundary
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11. 19.3 Standard Reduction Potentials
Designated Eo
Measured relative to the standard hydrogen electrode (SHE)
Standard conditions
Assigned a value of 0 V
Reaction
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Determine the overall cell reaction and
E°cell (at 25°C) of a galvanic cell made of
an Al electrode in a 1.0 M Al(NO3)3 solution
and a Cu electrode in a 1.0 M Cu(NO3)2
solution.
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Cell Potential
Anode
Cathode
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Cell Reaction
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21. 19.4 Spontaneity of Redox Reactions Under Standard-State Conditions
Based on electric charge and work
n is the number of moles of electrons
F is the Faraday constant
96,500 J/V . mol e
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At standard conditions
Relation to the equilibrium constant
Using the values of R, T (298 K) and F
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Determine the value of a) DGo and b) K for the following reaction.
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a) Standard free energy change, DGo
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b) Equilibrium constant, K
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19.5 Spontaneity of Redox Reactions Under Conditions Other than Standard-State Conditions
The Nernst Equation –derived from thermodynamics or
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Will the following reaction occur spontaneously at 298 K if [Fe2+] = M and [Cd2+] = M?
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Spontaneous since E is a positive even though very small.
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30. Concentration cells – galvanic cells composed of the same material but differing in ion concentrations 30

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19.6 Batteries
A battery is a galvanic cell, or a series of cells connected that can be used to deliver a self-contained source of direct electric current.
Dry Cells and Alkaline Batteries
no fluid components
Zn container in contact with MnO2 and an electrotyte
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Dry Cell
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Alkaline Cell
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Lead Storage Batteries
Six identical cells in series
Lead anode and PbO2 cathode
Immersed in H2SO4
Each cell delivers ~ 2 V
Rechargeable
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Lead Storage Battery
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Lithium-Ion Batteries
The overall cell potential is 3.4 V, which is a relatively large potential.
Lithium is also the lightest metal—only
6.941 g of Li (its molar mass) are needed to produce 1 mole of electrons.
recharged hundreds of times.
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Fuel Cells
Direct production of electricity by electrochemical means
Increased efficiency of power production
In its simplest form
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Simple Form of Fuel Cell
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19.7 Electrolysis
electrolysis - the use of electric energy to drive a nonspontaneous chemical reaction
electrolytic cell – the cell used to carry out electrolysis
same principles apply to both galvanic and electrolytic cells
in aqueous solutions you must also consider the oxidation or reduction of water
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Examples
Molten sodium chloride
Carried out in a Downs cell
Laboratory
Commercial
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Electrolysis of water
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Electrolysis of aqueous sodium chloride solution
Possible anode reactions
Possible cathode reactions
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Consider overvoltage – difference between estimated and actual voltage
Overvoltage for O2 is higher than Cl2
Cell reaction
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Quantitative Applications- measure the current (in amperes) that passes through an electrolytic cell in a given period of time.
Mass of product formed or reactant consumed
Proportional to the amount of electricity transferred
Molar mass
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Defining equation
Coulomb (C)
Second (s)
Ampere (A)
(coulomb/second)
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A constant current of A is passed
through an electrolytic cell containing molten
MgCl2 for 18 h. What mass of Mg is produced?
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19.8 Corrosion
Corrosion - generally refers to the deterioration of a metal by an electrochemical process.
Many metals undergo corrosion
For example, corrosion of Fe, oxidation of Al
Can be enhanced by atmospheric conditions ( e.g. acidic medium)
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Electrochemical Process for the Formation of Rust
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Electochemical processes can be used to prevent corrosion
Passivation – formation of a thin oxide layer by treating with an oxidizing agent
Formation of an alloy
Coating with a layer of a less active metal
Tin cans
Galvanization (zinc-plating)
Zinc oxide coating constitutes the protective coating
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Key Points
Balancing redox reactions
Half-reaction method
Galvanic cells
Electrodes (anode and cathode)
Cell potential (Ecell)
Cell diagrams
Standard Reduction potentials
Relative to the standard hydrogen electrode
Used to determine cell potential
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Spontaneity – standard conditions
Free energy
Equilibrium Constant
Spontaneity – nonstandard conditions
Nernst Equation
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Batteries
Dry cell and alkaline batteries
Lead storage battery
Lithium-ion batteries
Fuel cells
Electrolysis
Molten salts
Aqueous solutions
Quantitative applications
Corrosion
Metal deterioration
Prevention
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