Transcript ELECTROCHEMISTRY What is Electrochemistry  The branch of chemistry that deals with the use of spontaneous chemical reaction to produce electricity and the use.

ELECTROCHEMISTRY
What is Electrochemistry
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The branch of chemistry that deals with the use of
spontaneous chemical reaction to produce electricity
and the use of electricity to bring about non
spontaneous chemical change.
What are Half reactions?
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Oxidation occurs together with redox reaction but it is
helpful to consider each process separately by
writing two half reactions.
Mg(s) Mg2+(aq) + 2eFe3+(aq) + 3e-  Fe(s)
Class Practice
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Write the equation for the half reaction for (a) the
oxidation of iron (II) ions to iron(III) ions in aqueous
solution; (b) the reduction of copper (II) ions in
aqueous solution to copper metal.
Balancing redox reactions
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The chemical equation of a reduction half reaction
is added to that of an oxidation half reaction to
form the chemical equation for the overall redox
reaction.
You can balance a redox reaction in acidic and
basic reaction.
Class example
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Permanganate ions, MnO4-, react with oxalic acid,
H2C2O4, in acidic aqueous solution to produce
manganese(II)ions and carbon dioxide gas. The
partial skeletal equation is
MnO4- (aq) + H2C2O4(aq)  Mn2+(aq) +CO2(g)
In basic solution
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The products of the reaction of bromide ions with
permanganate ions in basic aqueous solution are
solid manganese (IV) oxide, MnO2, and bromate
ions, BrO3- .
Balance the chemical equation for the reaction.
Homework
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Galvanic cell
The Galvanic cell, named after Luigi Galvani, consists of two
different metals connected by a salt bridge or a porous disk
between the individual half-cells. It is also known as a voltaic
cell and an electrochemical cell.
What occurs in a galvanic cell:
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The zinc electrode is losing mass as Zn metal is oxidized to
Zn2+ ions which go into solution.
The concentration of the Zn2+ solution is increasing. Anions,
negative ions (e.g. SO42-), are flowing from the salt bridge
toward the anode to balance the positive charge of the Zn2+
ions produced.
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The copper electrode is gaining mass as Cu2+ ions in the solution are
reduced to Cu metal.
The concentration of the Cu2+ solution is decreasing.
Cations, positive ions (e.g. Na+), are flowing from the salt bridge toward
the cathode to replace the positive charge of the Cu2+ ions that consumed.
A reaction may start at standard-state conditions, but as the reaction
proceeds, the concentrations of the solutions change, the driving behind the
reaction becomes weaker, and the cell potential eventually reaches zero.
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In a galvanic cell, a spontaneous chemical reaction
tends to draw electrons into the cell through the
cathode, the site of reduction, and to release them
at the anode, the site of oxidation.
Notation For Cells
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Instead of drawing a cell diagram, chemists have
devised a shorthand way of completely describing
a cell called line notation. This notation scheme
places the constituents of the cathode on the right
and the anode components on the left. For example,
a half-cell containing 1M solutions of CuO and HCl
and a Pt electrode for the reduction of Cu2+ would
be written as: Pt (s) | Cu2+ (aq), H+ (aq)
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Note that the spectator ions, oxide and chloride,
have been left out of the notation and that the
anode will be written to the far left.
The salt bridge or porous disk is shown in the
notation as a double line ( || ). Therefore, a cell
that undergoes the oxidation of magnesium by Al3+
could have the following cell notation if the anode is
magnesium and the cathode is aluminum:
Mg (s) | Mg2+ (aq) || Al3+ (aq) | Al (s)
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An electrode is designated by representing the
interfaces between phases by Ι.
A cell diagram depicts the physical arrangement of
species and inter faces with a salt bridge denoted
by II.
Class Practice
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Write the diagram for a cell that has a hydrogen
electrode on the left, an iron (III)/iron(II) electrode
on the right, and includes a salt bridge. Both
electrode contacts are platinum.
Cell Potential
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An idealized cell for the electrolysis of sodium chloride is shown in the figure below.
A source of direct current is connected to a pair of inert electrodes immersed in
molten sodium chloride. Because the salt has been heated until it melts, the Na+ ions
flow toward the negative electrode and the Cl- ions flow toward the positive
electrode.
When Na+ ions collide with the negative electrode, the battery carries a large enough
potential to force these ions to pick up electrons to form sodium metal.
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Negative electrode (cathode): Na+ + e--  Na
Cl- ions that collide with the positive electrode are oxidized to
Cl2 gas, which bubbles off at this electrode.
Positive electrode (anode): 2 Cl-  Cl2 + 2 eThe net effect of passing an electric current through the molten
salt in this cell is to decompose sodium chloride into its
elements, sodium metal and chlorine gas.
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Electrolysis of NaCl:
Cathode (-): Na+ + e- Na
Anode (+): 2 Cl- Cl2 + 2 eThe potential required to oxidize Cl- ions to Cl2 is -1.36 volts
and the potential needed to reduce Na+ ions to sodium
metal is -2.71 volts. The battery used to drive this reaction
must therefore have a potential of at least 4.07 volts.
This example explains why the process is called electrolysis.
Electrolysis uses an electric current to split a compound into
its elements.
Electrolysis: 2 NaCl(l) 2 Na(l) + Cl2(g)
Home work
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Standard cell potential
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A cell's standard state potential is the potential of the cell
under standard state conditions, which is approximated
with concentrations of 1 mole per liter (1 M) and
pressures of 1 atmosphere at 25oC.
To calculate the standard cell potential for a reaction:
Write the oxidation and reduction half-reactions for the
cell.
Look up the reduction potential, Eoreduction, for the
reduction half-reaction in a table of reduction potentials
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Look up the reduction potential for the reverse of the
oxidation half-reaction and reverse the sign to obtain the
oxidation potential. For the oxidation half-reaction,
Eooxidation = - Eoreduction.
Add the potentials of the half-cells to get the overall
standard cell potential.
Eocell = Eoreduction + Eooxidation
Example:
Find the standard cell potential for an electrochemical cell
with the following cell reaction.
Zn(s) + Cu2+(aq) Zn2+(aq) + Cu(s)
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Write the half-reactions for each process.
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Zn(s)  Zn2+(aq) + 2 e- Cu2+(aq) + 2 e-  Cu(s)
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Look up the standard potentials for the reduction half-reaction.
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Eoreduction of Cu2+ = + 0.339 V
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Look up the standard reduction potential for the reverse of the oxidation
reaction and change the sign.
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Eoreduction of Zn2+ = - 0.762 V Eooxidation of Zn = - ( - 0.762 V) = + 0.762 V
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Add the cell potentials together to get the overall standard cell potential.
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oxidation: Zn(s)  Zn2+(aq) + 2 e- Eoox. = - Eored. = - (- 0.762 V) =
+ 0.762 V
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reduction: Cu2+(aq) + 2 e-  Cu(s) Eored. = + 0.339 V overall: Zn(s) +
Cu2+(aq)  Zn2+(aq) + Cu(s) Eocell = + 1.101 V
Class Practice
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The standard potential of a Zn2+/Zn electrode is –
0.76 V, and the standard potential of the cell
Zn(s)IZn2+(aq)IICu2+(aq)ICu(s) is 1.10 V. What is the
standard potential of the Cu2+/Cu electrode?
Home work
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The Electrochemical Series
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The oxidizing and reducing power of a substance
can be determined by its position in the
electrochemical series. The strongest oxidizing
agents are at the top of the table as reactants, the
strongest reducing agents are at the top of the
table as products.
Class practice
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Can aqueous potassium permanganate be used to
oxidize iron(II) to iron(III) under standard conditions
in acidic solution.
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To find a standard cell potential arising from a
spontaneous reaction we must combine the standard
potential of the cathode half reaction(reduction) with
that of the anode half reaction (oxidation) in such a
way so as to obtain a positive value. The overall
potential must be positive because that corresponds to
a spontaneous process, and only a spontaneous process
can generate a potential. If the calculation results in a
negative value, this means that the reverse reaction is
spontaneous.
Homework
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Standard potential, free energy and
equilibrium constant
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The free energy change is a measure of the change
in the total entropy of a system and its surroundings
at constant pressure; spontaneous processes are
accompanied by a decrease in free energy.
G=H−TS
H is the enthalpy
T is the temperature
S is the entropy
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When the reaction free energy ∆Gf is negative ,
the cell reaction is spontaneous and the cell
generates a positive potential. When ∆Gf is is
large as well as negative, the cell potential is high
as well as positive. The relationship suggests that
∆Gf =−nFE
n is the number of moles of electrons that are
transferred between the electrodes for the cell
reaction as written in the chemical equation.
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The Faraday constant, F is the magnitude of the
charge per mole of electrons: F=Nae=9.6485 X104
C/mol.
1C.V=1J
We can write F= 9.6485 X104 J/V.mol
Class Practice
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The cell Cr(s) ΙCr2+(aq)Ι Cu (s) was found to have E⁰
= +1.08V at 298K.(a) Write the balanced net
equation for the cell reaction;(b) determine n;
and© calculate the standard reaction free energy
at 298K.
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The following cell was set up: Hg(l) ΙHg2Cl₂(s)ΙHCl
(aq)ΙΙHg₂(NO₃)₂(aq)ΙHg(l), E⁰ =+ 0.52V at 298.(a)
Write the equation for the cell reaction.(b)
determine n, and © calculate the standard
reaction free energy at 298K.
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The equilibrium constant of a reaction can be
calculated from standard potentials by combining
the equation for the half reactions to give the
reaction of interest and determining the standard
potential of the corresponding cell.
Class Practice
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The reaction between zinc metal and iodine in
water generates 1.30 v under standard conditions.
Determine a)n and b) ∆Gf ⁰ for the cell reaction
Zn(s)+I2(aq) Zn2+ (aq) + 2I- (aq)
The solubility product is the equilibrium constant
for the dissolution of a salt. Calculate the solubility
product of silver chloride.
The Nernst Equation
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The variation of cell potential with composition is
expressed by the Nernst equation:
E=E⁰ ─(RT/nF) ln Q
Class Practice
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Calculate the potential at 25⁰C of a Daniel cell in
which the concentration of Zn²+ ions is 0.10 mol/L
and that of the Cu2+ ions is 0.0010 mol/L
Home work
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Electrolysis
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Electrolysis is a method of separating chemically
bonded elements and compounds by passing an
electric current through them.
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In an electrolytic cell, current supplied by an
external source is used to drive a nonspontaneous
redox reaction.
The Galvanic cell, consists of two different metals
connected by a salt bridge or a porous disk
between the individual half-cells. It is also known as
a voltaic cell and an electrochemical cell.
The porous bridge is substituted by a
Salt bridge for a galvanic cell.
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Oxidation-reduction or redox reactions take place in
electrochemical cells. There are two types of
electrochemical cells. Spontaneous reactions occur in
galvanic (voltaic) cells; nonspontaneous reactions occur
in electrolytic cells.
Both types of cells contain electrodes where the
oxidation and reduction reactions occur. Oxidation
occurs at the electrode termed the anode and reduction
occurs at the electrode called the cathode.
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The anode of an electrolytic cell is positive (cathode
is negative), since the anode attracts anions from
the solution. However, the anode of a galvanic cell
is negatively charged, since the spontaneous
oxidation at the anode is the source of the cell's
electrons or negative charge. The cathode of a
galvanic cell is its positive terminal. In both galvanic
and electrolytic cells, oxidation takes place at the
anode and electrons flow from the anode to the
cathode.
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The redox reaction in a galvanic cell is a
spontaneous reaction. Hence, galvanic cells are
commonly used as batteries. Galvanic cell reactions
supply energy which is used to perform work. The
energy is harnessed by situating the oxidation and
reduction reactions in separate containers, joined by
an apparatus that allows electrons to flow. A
common galvanic cell is the Daniell cell.
Class practice
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Predict the products resulting from the electrolysis of
1M ZnNO2(aq) at pH=7
Home work
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