Transcript Electrochemistry

Electrochemistry
The study of the interchange of
chemical and electrical energy.
Electrochemical Reactions
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All electrochemical reactions involve the transfer
of electrons and are therefore, oxidationreduction reaction.
Electrons are transferred from the reducing
agent to the oxidizing agent.
Oxidation is a loss of electrons (increase in
oxidation number)-”OIL”
Reduction is a gain of electrons (decrease in
oxidation number)-”RIG”
Types of Electrochemical Cells
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Galvanic or Voltaic Cells-those in which a
spontaneous chemical reaction produces
an electrical current that can be used to
do work.
Electrolytic Cells-those in which electrical
energy from an outside source causes a
nonspontaneous reaction to occur.
Components of a Galvanic Cell
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Cell-the reacting system
Electrodes-surfaces where the electric current exits or
enters
1) anode- electrode compartment in which oxidation
occurs. “AN OX”
2) cathode-electrode compartment in which reduction
occurs. “RED CAT”
Salt Bridge- U-tube filled with an electrolyte or a porous
disk in a tube connecting the two solutions.
Wire-path by which the electrons flow from one
compartment to the other.
Electrons flow through the wire from the reducing agent
to the oxidizing agent (from the anode to the cathode)
Diagram of Galvanic Cell
Cell Potential
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Cell potential (Ecell) or electromotive force
(emf) is the “pull” or driving force on the
electrons.
The unit of electrical potential is the volt
(V) which is defined as 1 joule/coulomb.
Cell potential is measured with a
voltmeter.
Standard Reduction Potentials
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Reactions in galvanic cells are broken down into halfreactions with each being assigned a reduction potential.
All half reactions are assigned reduction potentials using
the standard hydrogen electrode as the reference. (see
page 796)
The potentials are all given as reduction processes.
If the process must be reversed (oxidation process), the
sign for the potential is reversed.
Since reduction potential is an intensive process (doesn’t
depend on the how many times the reaction occurs), the
value of the reduction potential is not changed when a
half-reaction is multiplied by an integer to balance an
equation.
Standard Reduction Potentials
(continued)
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The more positive the Eo value for a halfreaction, the greater tendency for the halfreaction to occur.
The more negative the Eo value for a halfreaction, the greater tendency for the halfreaction to occur in the opposite direction.
If Eocell > 0 (positive), the forward reaction is
spontaneous.
If Eocell < 0 (negative), the forward reaction is
not spontaneous and would have to be carried
out in an electrolytic cell.
Complete the practice problems
on page 797.
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A. 0.71 V
B. 0.32 V
Zinc-Copper Galvanic Cell
The Zinc-Copper Cell
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Example:
Zn(s) + Cu2+(aq)  Zn2+ (aq) + Cu(s)
Anode: Zn  Zn2+ + 2 eCathode: Cu2+ + 2e-  Cu
Eocell = .337 + .763 = 1.10 V
Line Notation:
Zn | Zn2+ | | Cu2+ | Cu (anode is written on the
left side and the vertical line represents a phase
difference or boundary)
Write the line notation for a
galvanic cell consisting of
copper (II) and chromium (III)
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Cr3+ + 3e-  Cr
Eo = -0.74 V
Cu2+ + 2e-  Cu Eo = 0.337 V
Copper reduction occurs at the cathode
Chromium oxidation occurs at the anode.
Line Notation:
Cr | Cr3+ | | Cu2+ | Cu
Cell Potential, Work, and Free Energy
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The work that can be accomplished when
electrons are transferred through a wire
depends on the “push” behind the electrons.
Potential difference (V) = work (J)/charge (C)
E = -w/q
Work is viewed from the point of view of the
system. (Work flowing out of a system is
indicated by a minus sign).
In any real, spontaneous process some energy is
wasted due to frictional heating-the actual work
realized is always less than the calculated
maximum.
Electrical Charge
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The charge on one mole of electrons is a
constant called the faraday (F), which has
the value 96,485 coulombs of charge per
mole of electrons.
q = nF (n is the number of moles of e-)
w (∆G) = -nFEmax
Solve example 17-3 on page 802.
The Nernst Equation
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The Nernst Equation is used to calculate
electrode potential and cell potentials for
concentrations and partial pressures other than
standard-state values.
E = Eo – (2.303 RT/nF ) log Q
E = potential under nonstandard conditions
Eo = standard potential
R= 8.314
T = temp in Kelvin
n= number of moles of electrons transferred
F = 96,485 C/mole of eQ=reaction quotient
Electrolytic Cells
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In an electrolytic cell, an outside source of
voltage is used to force a nonspontaneous redox
reaction to take place.
Oxidation takes place at the anode and
reduction takes place at the cathode just as it
does in a galvanic cell.
The cell potential in an electrolytic cell < 0.
Electrolytic cells are used in electroplating.
Electrolysis Problems
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I = q/t
I = current (amperes, A)
1 amp = 1C/sec
q = charge (coulombs, C)
t = time (sec)
Once the charge is known, solve the
problem as a stoichiometry problem.
Practice Problem #1
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How long must a current of 5.00 A be
applied to a solution of Ag+ to produce
10.5 g of silver metal?
Practice Problem #2
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What mass of Co can be produced from
aqueous Co2+ in 1 hour with a current of
15 A?
Practice Problem #3
A zinc-copper battery is constructed as follows at
25oC:
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Zn2+ (0.1M)
Cu2+ (2.50M) Cu
The mass of each electrode is 200g.
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Calculate the cell potential when this battery is first
connected.
Calculate the cell potential after 10.0A of current has
flowed for 10.0 h.
Calculate the mass of each electrode after 10.0 h
How long can this battery deliver a current of 10.0A before
it goes dead?