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Tro, Chemistry: A Molecular Approach 1
Electrochemistry
19.1 Redox Reactions.
19.2 Galvanic Cells.
19.3 Standard Reaction Potential.
19.4 Spontaneity of Redox Reactions.
19.5 The Effect of Concentration on Cell Emf.
19.6 Batteries.
19.7 Corrosion.
19.8 Electrolysis.
3
Terms used with electricity:
•Galvanometer - an instrument for detecting electric current.
•Voltmeter - an instrument to measure the potential difference between two
half-cells in a voltaic cell.
•Amperes - the SI units for current flow = # e- / s
•Volts - a measurement of electric potential difference between two
electrodes.
•Electromotive force, emf - another term for volts.
•Cations - positive ions attracted to the cathode.
•Anions - negative ions attracted to the anode.
•Electrolysis - the process by which an electric current produces a chemical
change.
•Voltaic cell - a device used to produce electric energy from an oxidation-
reduction reaction.
•Battery - two or more electrochemical cells operating as a unit.
•External circuit - a "wires" connected to a battery providing a path for
electricity to flow.
•Internal circuit - the electrolyte inside a battery through which ions can
move.
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19.1 Redox Reactions
Electrochemistry is the branch of chemistry that deals
with the interconversion of electrical energy and
chemical energy.
Electrochemical processes are redox (oxidation -
reduction) reactions.
In redox reactions, electrons are transferred from one
substance to another.
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Redox Reaction• one or more elements change oxidation number
all single displacement, and combustion,
some synthesis and decomposition
• always have both oxidation and reduction
split reaction into oxidation half-reaction and a
reduction half-reaction
• aka electron transfer reactions
half-reactions include electrons
• oxidizing agent is reactant molecule that causes oxidation
contains element reduced
• reducing agent is reactant molecule that causes reduction
contains the element oxidized
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Oxidation & Reduction
• oxidation is the process that occurs whenelement loses electrons
compound adds oxygen
oxidation number of an element increases
compound loses hydrogen
half-reaction has electrons as products
• reduction is the process that occurs whenoxidation number of an element decreases
element gains electrons
compound loses oxygen
compound gains hydrogen
half-reactions have electrons as reactants
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Rules for Assigning Oxidation States
5. in their compounds, nonmetals have oxidation
states according to the table below
nonmetals higher on the table take priority
Nonmetal Oxidation State Example
F -1 CF4
H +1 CH4
O -2 CO2
Group 7A -1 CCl4
Group 6A -2 CS2
Group 5A -3 NH3
Balancing Redox Equations
Suppose we are asked to balance the equation showing
the oxidation of Fe(III) ions by dichromate ions () in
acidic medium.
The following steps will help us balance the equation
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Practice - Balance the Equation
ClO3-1 + Cl-1 Cl2 (in acid)
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Practice - Balance the Equation
ClO3-1 + Cl-1 Cl2 (in acid)
+5 -2 -1 0
oxidationreduction
ox: 2 Cl-1 Cl2 + 2 e-1 } x5
red: 2 ClO3-1 + 10 e-1 + 12 H+ Cl2 + 6 H2O} x1
tot 10 Cl-1 + 2 ClO3-1 + 12 H+ 6 Cl2 + 6 H2O
1 ClO3-1 + 5 Cl-1 + 6 H+1 3 Cl2 + 3 H2O
Worked Example 19.1
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Electrochemical Cells• electrochemistry is the study of redox reactions
that produce or require an electric current
• the conversion between chemical energy and electrical energy is carried out in an electrochemical cell
• spontaneous redox reactions take place in a voltaic cell
aka galvanic cells
• nonspontaneous redox reactions can be made to occur in an electrolytic cell by the addition of electrical energy
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Electrochemical Cells• oxidation and reduction reactions kept separate
half-cells
• electron flow through a wire along with ion flow
through a solution constitutes an electric circuit
• requires a conductive solid (metal or graphite)
electrode to allow the transfer of electrons
through external circuit
• ion exchange between the two halves of the system
electrolyte
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Electrodes• Anodeelectrode where oxidation occursanions attracted to itconnected to positive end of battery in electrolytic
cellloses weight in electrolytic cell
• Cathodeelectrode where reduction occurscations attracted to itconnected to negative end of battery in electrolytic
cellgains weight in electrolytic cellelectrode where plating takes place in electroplating
19.2 Galvanic Cell
The experimental apparatus for generating
electricity through the use of a spontaneous
reaction is called a galvanic cell or voltaic cell,
Figure 19.1 shows the essential components of a
galvanic cell.
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The difference in electrical potential between the anode
and the cathode is measured by a voltameter (Figure 19.2)
The voltage across the electrodes of a galvanic cell is
called the cell voltage, or cell potential .
Another common term for cell voltage is the
electromotive force or emf (E), which is a measure of
voltage, not force.
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Current and Voltage• the number of electrons that flow through the system per
second is the current unit = Ampere
1 A of current = 1 Coulomb of charge flowing by each second
1 A = 6.242 x 1018 electrons/second
Electrode surface area dictates the number of electrons that can flow
• the difference in potential energy between the reactants and products is the potential difference unit = Volt
1 V of force = 1 J of energy/Coulomb of charge
the voltage needed to drive electrons through the external circuit
amount of force pushing the electrons through the wire is called the electromotive force, emf
Tro, Chemistry: A Molecular Approach 29
Cell Potential• the difference in potential energy between the
anode and the cathode in a voltaic cell is called the cell potential
• the cell potential depends on the relative case with which the oxidizing agent is reduced at the cathode and the reducing agent is oxidized at the anode
• the cell potential under standard conditions is called the standard emf, E°cell
25°C, 1 atm for gases, 1 M concentration of solution
sum of the cell potentials for the half-reactions
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Standard Reduction Potential• a half-reaction with a strong tendency to
occur has a large + half-cell potential
• when two half-cells are connected, the electrons will flow so that the half-reaction with the stronger tendency will occur
• we cannot measure the absolute tendency of a half-reaction, we can only measure it relative to another half-reaction
• we select as a standard half-reaction the reduction of H+ to H2 under standard conditions, which we assign a potential difference = 0 v standard hydrogen electrode, SHE
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19.3 Standard Reduction Potentials
19.3
Standard reduction potential (E0) is the voltage associated with
a reduction reaction at an electrode when all solutes are 1 M and
all gases are at 1 atm.
E0 = 0 V
Standard hydrogen electrode (SHE)
Reduction Reaction
2e- + 2H+ (1 M) H2 (1 atm)
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Half-Cell Potentials• SHE reduction potential is defined to be exactly 0 v
• half-reactions with a stronger tendency toward
reduction than the SHE have a + value for E°red
• half-reactions with a stronger tendency toward
oxidation than the SHE have a value for E°red
• E°cell = E°oxidation + E°reduction
E°oxidation = E°reduction
when adding E° values for the half-cells, do not multiply the
half-cell E° values, even if you need to multiply the half-
reactions to balance the equation
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Worked Example 19.2
Worked Example 19.3
19.4 Spontaneity of Redox Reaction
How E⁰ for cell is related to thermodynamic
quantities such as ΔG⁰ and K
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E°cell, DG° and K
• for a spontaneous reaction
one the proceeds in the forward direction with the
chemicals in their standard states
DG° < 1 (negative)
E° > 1 (positive)
K > 1
• DG° = −RTlnK = −nFE°cell
n is the number of electrons
F = Faraday’s Constant = 96,485 C/mol e−
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Worked Example 19.4
Worked Example 19.5
19.5 The Effect of Concentration on Cell Emf
There is a mathematical relationship between
the emf of a galvanic cell and the
concentration of reactants and products in a
redox reaction under nonstandard-state
conditions. This equation called
Nernest Equation
E = E⁰ - RT /nF ln [oxd]/[red]
From equation (18.13) page 821
ΔG = ΔG⁰ + RT ln Q
because : ΔG = -nFE and ΔG⁰ = -nFE⁰
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During the operation of a galvanic cell, electrons flow
from the anode to the cathode, resulting in product
formation and a decrease in reactant concentration.
Thus Q increases, which means that E decreases, the cell
reaches equilibrium.
At equilibrium, there is no net transfer of electrons,
so E= 0 and Q = K, where K is the equilibrium constant
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Worked Example 19.6
Worked Example 19.7
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Concentration Cells• it is possible to get a spontaneous reaction when the oxidation
and reduction reactions are the same, as long as the electrolyte concentrations are different
• the difference in energy is due to the entropic difference in the solutions
the more concentrated solution has lower entropy than the less concentrated
• electrons will flow from the electrode in the less concentrated solution to the electrode in the more concentrated solution
oxidation of the electrode in the less concentrated solution will increase the ion concentration in the solution – the less concentrated solution has the anode
reduction of the solution ions at the electrode in the more concentrated solution reduces the ion concentration – the more concentrated solution has the cathode
Concentration Cell
Because electrode potential depends on ion concentration, it is possible to construct a
galvanic cell from two half-cells composed of
the same material but differing in ion
concentrations. Such a cell is called
a concentration cell.
Tro, Chemistry: A Molecular Approach 56
Concentration Cell
when the cell concentrations
are different, electrons flow
from the side with the less
concentrated solution
(anode) to the side with the
more concentrated solution
(cathode)
Cu(s) Cu2+(aq) (0.010 M) Cu2+
(aq) (2.0 M) Cu(s)
19.6 Batteries
A battery is a galvanic cell, or a series of combined
galvanic cells, that can be used as a source of direct
electric current at a constant voltage.
The operation of battery is similar in principle to that
of the galvanic cells.
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19.7 Corrosion
Corrosion is the term usually applied to the deterioration
of metals by an electrochemical process. There are many
examples of corrosion: Rust on iron, tarnish of silver, and the
green patina formed on copper and brass.
Corrosion causes damage to buildings, bridges, ships, and cars.
The most familiar example of corrosion is the formation of rust
on iron.
Oxygen gas and water must be present for iron to rust.
Corrosion
The main steps are believed to be as follows:
A region of the metal’s surface serves as the anode,
The oxedation reaction is:
Fe(s) Fe²+(aq) + 2e
The electron given up by iron reduce atmospheric oxygen
to water at the cathode, which is another region of the
same metal surface:
O2 + 4H + 4e 2H2O(l)
The overall reaction is:
2Fe(s) + O2(g) + 4H(aq) 2Fe²+(aq) + 2H2O(l)
We find the standard emf for this process:
Eºcell = Eºcathode – Eºanode
=1.23 V – (-0.44 V) = 1.67 V
Note that this reaction occurs in an acidic
medium, see figure 19.14 which shows the
mechanism of the rust formation.
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Tro, Chemistry: A Molecular Approach 66
Cathodic protection is a process in which the metal that is
to be protected from corrosion is made the cathodic in
what amounts to a galvanic cell.
Figure 19.15 shows how an iron nail can be protected
from rusting by connecting the nail to a piece of zinc.
With out such protection, an iron nail quickly rusts in
water.
Rusting of underground iron pipes and iron storage tanks
can be prevented or greatly reduced by connecting them
to metals such as zinc and magnesium, which oxidize
more readily than iron as in figure 19.16
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Figure 19.15
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19.8 Electrolysis
In contrast to spontaneou redox reactions, which result
in the conversion of chemical energy into electrical
energy, electrolysis is the process in which electrical
energy is used to cause a nonspontaneous chemical
reaction to occure.
An electrolytic cell is an apparatus for carrying out
electrolysis.
Electrolysis of molten Sodium Chloride
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Faraday’s Law
• the amount of metal deposited during
electrolysis is directly proportional to the charge
on the cation, the current, and the length of time
the cell runs
charge that flows through the cell = current x time
Quantitative Aspects of Elctrolysis
The quantitative treatment of electrolysis was developed primarily
by Faraday.
He observed that:
The mass of product formed (or reactant consumed) at an
electrode is proportional to both the amount of electricity
transferred at the electrode and the molar mass of the substance.
The relationship between charge (in coulombs, C) and current is:
1C = 1A x 1s
That is a coulomb is the quantity of electrical charge passing any
point in the circuit in 1 second when the current is 1 ampere.
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Figure 19.20 shows the steps involved in calculating the quantities
of substances produced in electrolysis.
Finish
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Terms used with electricity:
•Galvanometer - an instrument for detecting electric current.
•Voltmeter - an instrument to measure the potential difference between two half-
cells in a voltaic cell.
•Amperes - the SI units for current flow = # e- / s
•Volts - a measurement of electric potential difference between two electrodes.
•Electromotive force, emf - another term for volts.
•Cations - positive ions attracted to the cathode.
•Anions - negative ions attracted to the anode.
•Electrolysis - the process by which an electric current produces a chemical change.
•Voltaic cell - a device used to produce electric energy from an oxidation-reduction
reaction.
•Battery - two or more electrochemical cells operating as a unit.
•External circuit - a "wires" connected to a battery providing a path for electricity to
flow.
•Internal circuit - the electrolyte inside a battery through which ions can move.