Introduction to Electrochemistry

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Introduction to Electrochemistry
Read Chapter 14.
Electrochemistry can be broadly defined as the
study of charge-transfer phenomena. As such, the
field of electrochemistry includes a wide range
of different chemical and physical phenomena.
These areas include (but are not limited to)
battery chemistry, photosynthesis, ion-selective
electrodes, coulometry, and many biochemical
processes. Although wide ranging,
electrochemistry has found many practical
applications in analytical measurements.
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Electro-analytical Chemistry
Electro-analytical chemistry is the field of
electrochemistry that utilizes the relationship
between chemical phenomena which involve charge
transfer (e.g. redox reactions, ion separation,
etc.) and the electrical properties that
accompany these phenomena for some analytical
determination. This relationship is further
broken down into fields based on the type of
measurement that is made. Potentiometry involves
the measurement of potential for quantitative
analysis, and electrolytic electrochemical
phenomena involve the application of a potential
or current to drive a chemical phenomenon,
resulting in some measurable signal which may be
used in an analytical determination.
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There are two parts to understanding
electrochemistry the first is thermodynamics
the second is the kinetics of electrode
processes. For the latter, one needs to study
the surface chemistry to obtain a real
understanding of how electrochemical systems
work. For chemists to understand the principles
underlying the functioning of such practically
important systems as batteries, fuel cells,
corrosion, electrolysis, as well as membranes and
biomembranes (of utmost importance for the
understanding of, drug delivery and the
functioning of cell membranes) they must
therefore be taught the basics of interfacial
structure, electrochemical kinetics and transport
processes. (From International Society of
Electrochemistry) site
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Example of a Galvanic Cell
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Redox reactions involve electron
transfer. Acid-Base reactions involve proton
transfer. The key difference is that electrons
can be transported through space via wires or any
other conducting device with the result that the
oxidation and reduction reactions can occur in
different places. Protons are transported in an
aqueous (or some other polar) environment, so
acid-base reactions occur in the same place.
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Basic Concepts
Redox reactions involve a species which is
oxidized and another that is reduced.
In the above, Fe3 is reduced to Fe2 . It is
the oxidizing agent. Since DGlt0 for this
reaction we can say that V3 wants the extra
electron less than Fe3.
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Galvanic Cell
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An Aside Why wont this cell work?
Ag will go to left electrode and ask for e from
Cd(s) directly.
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Will this cell work?
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How badly to the electrons want to flow?
I current in amps R resistance in ohms E
potential difference in Volts
q n x F
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Voltaic Cells Electrochemical cells that use an
oxidation-reduction reaction to generate an
electric current are known as galvanic or voltaic
cells. Let's take another look at the voltaic
cell in the figure below.
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Voltaic Cells
Taken from http//chemed.chem.purdue.edu/genchem
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Within each half-cell, reaction occurs on the
surface of the metal electrode. At the zinc
electrode, zinc atoms are oxidized to form Zn2
ions, which go into solution. The electrons
liberated in this reaction flow through the zinc
metal until they reach the wire that connects the
zinc electrode to the platinum wire. They then
flow through the platinum wire, where they
eventually reduce an H ion in the neighboring
solution to a hydrogen atom, which combines with
another hydrogen atom to form an H2 molecule.
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The electrode at which oxidation takes place in a
electrochemical cell is called the anode. The
electrode at which reduction occurs is called the
cathode. The identity of the cathode and anode
can be remembered by recognizing that positive
ions, or cations, flow toward the cathode, while
negative ions, or anions, flow toward the anode.
In the voltaic cell shown above, H ions flow
toward the cathode, where they are reduced to H2
gas. On the other side of the cell, Cl- ions are
released from the salt bridge and flow toward the
anode, where the zinc metal is oxidized.
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The voltaic cell consist of the two reactions.
oxidation
reduction

Or equivalently we can write the reactions as
follows
We can only measure E for the full reaction. We
would like to calculate E for the half
reactions. Before doing this, we must recognize
the E depends on concentrations.
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Voltaic Cells
Taken from http//chemed.chem.purdue.edu/genchem
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Since reactants and products are in their
standard states, we call the E for this cell the
standard reduction potential (Eo). Here Eo .76V.

We arbitrarily define the potential for, one half
reaction, the second reaction above to be exactly
0V when reactants and products are in their
standard states.
Since Eo for the cell is the sum of Eos for the
two half reactions we see that Eo for the first
half reaction is .76V.
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Oxidizing Power Increases
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Voltaic Cells
Taken from http//chemed.chem.purdue.edu/genchem
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This voltaic cell on the previous slide is fully
described with the following notation
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Line Notation For Voltaic Cells

• Voltaic cells can be described by a line notation
  based on the following conventions.
• Single vertical line indicates change in state or
  phase.
• Within a half-cell, the reactants are listed
  before the products.
• Activities of aqueous solns are written in
  parentheses after the symbol for the ion or
  molecule.
• A double vertical line indicates a junction
  between half-cells.
• The line notation for the anode (oxidation) is
  written before the line notation for the cathode
  (reduction).
• The line notation for a standard-state Daniell
  cell is written as follows.

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Electrons flow from the anode to the cathode in a
voltaic cell. (They flow from the electrode at
which they are given off to the electrode at
which they are consumed.) Reading from left to
right, this line notation therefore corresponds
to the direction in which electrons flow.
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The Nernst Equation (14-4 of book)
The Nernst equation relates the potential of a
cell in its standard state to that of a cell not
in its standard state. Consider the reaction
below.
We know from Le Chateliers principle that
increasing the concentration of Zn2 should drive
the reaction to the right. In other words it
should decrease the potential of the half cell.
The Nernst equation allows us to calculate this
increase for the above half reaction as
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The Nernst Equation Continued
The Nernst Eq. for the reaction
is
At 25oC this equation simplifies to
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The Nernst Equation For Complete Cell
Here E and E- are the potentials of the half
cells connected to the positive and negative
terminals of potentiometer respectively. Lets
consider an example.
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Voltaic Cells
-

Taken from http//chemed.chem.purdue.edu/genchem
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E and E- are potentials of half cells connected
to positive and negative terminals of
potentiometer respectively
See page 295 for another example
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Calculating Equilibrium Constants
Consider the reaction
made up of the following two half reactions
Eo1.700V Eo0.767V
Since Eo is greater for cerium this reaction will
be the reduction reaction. The standard
potential for the galvanic cell would be
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Calculating Equilibrium Constants Continued
Consider the reaction
In a galvanic cell we would have
At equilibrium E0 and
This connection to free energy is important
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Calculating Equilibrium Constants for Nonredox
Reactions (14-5)
Consider the reaction
This is a Ksp problem. Not a redox problem.
Nonetheless we can use electrochemistry to
calculate Ksp by considering
(at 25oC)
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Latimer Diagrams (Box 14-2 of book)
The oxidation states of elements are related to
each other
(7) (5) (1)
(0) (-1)
1.299
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Electrochemistry Skills

• Understand how voltaic cells work.
• Be able to calculate standard reduction
  potentials for voltaic cells, given the chemical
  reactions.
• Be able to describe a voltaic cell using the line
  notation and visa versa. Know which way
  electrons flow and where the anode and cathode
  are.
• Know how to work with the Nernst Eq. to include
  concentration dependencies and calculate
  equilibirum constants.
• Know the relation between E and DG and can use
  this relation in constructing Latimer diagrams.