Electrochemistry (download)

forestercuckooΜηχανική

27 Οκτ 2013 (πριν από 3 χρόνια και 7 μήνες)

77 εμφανίσεις

Electrochemistry

Thermodynamics at the electrode

Learning objectives


You will be able to:


Identify main components of an electrochemical cell


Write shorthand description of electrochemical cell


Calculate cell voltage using standard reduction potentials


Apply Nernst equation to determine free energy change


Apply Nernst equation to determine pH


Calculate K from electrode potentials


Calculate amount of material deposited in electrolysis





Energy in or energy out


Galvanic

(or
voltaic
) cell relies on
spontaneous process to generate a potential
capable of performing work


energy out


Electrolytic

cell performs chemical reactions
through application of a potential


energy in

Redox Review


Oxidation is...


Loss of electrons


Reduction is...


Gain of electrons


Oxidizing agents oxidize and are reduced


Reducing agents reduce and are oxidized

Redox at the heart of the matter


Zn displaces Cu from CuSO
4
(aq)


In direct contact the enthalpy of reaction is
dispersed as heat, and no useful work is done


Redox process:


Zn is the reducing agent


Cu
2+

is the oxidizing agent


Separating the combatants


Each metal in touch with a solution of its own ions


External circuit carries electrons transferred during the redox process


A “salt bridge” containing neutral ions completes the internal circuit.


With no current flowing, a potential develops


the potential for work


Unlike the reaction in the beaker, the energy released by the reaction
in the cell can perform useful work


like lighting a bulb

Labelling the parts

Odes to a galvanic cell


Cathode


Where reduction occurs


Where electrons are
consumed


Where positive ions
migrate to


Has positive sign



Anode


Where oxidation occurs


Where electrons are
generated


Where negative ions
migrate to


Has negative sign

The role of inert electrodes


Not all cells start with elements as the redox
agents


Consider the cell



Fe can be the anode but Fe
3+

cannot be the
cathode.


Use the Fe
3+

ions in solution as the
“cathode” with an inert metal such as Pt

Anode

Catho
de

Oxidati
on

Reduct
ion

Cell notation


Anode on left, cathode on right


Electrons flow from left to right


Oxidation on left, reduction on right


Single vertical = electrode/electrolyte boundary


Double vertical = salt bridge

Anode:

Zn
→Zn
2+

+ 2e

Cathode:

Cu
2+

+ 2e
→Cu

Vertical │denotes different phase


Fe(s)
│Fe
2+
(aq)║Fe
3+
(aq),Fe
2+
(aq)│Pt(s)



Cu(s)
│Cu
2+
(aq)║Cl
2
(g)│Cl
-
(aq)│C(s)

Connections: cell potential and free
energy


The cell in open circuit generates an
electromotive force (emf) or potential or
voltage. This is the potential to perform
work


Energy is charge moving under applied
voltage

Relating free energy and cell
potential


The Faraday:

F = 96 485 C/mol e



Standard conditions (1 M, 1 atm, 25
°
C)

Standard Reduction Potentials


The total cell potential is the sum of the potentials
for the two half reactions at each electrode


E
cell

= E
cath

+ E
an


From the cell voltage we cannot determine the
values of either


we must know one to get the
other


Enter the
standard hydrogen electrode (SHE)


All potentials are referenced to the SHE (=0 V)

Unpacking the SHE


The SHE consists of a Pt electrode in contact with
H
2
(g) at 1 atm in a solution of 1 M H
+
(aq).


The voltage of this half
-
cell is defined to be 0 V


An experimental cell containing the SHE half
-
cell
with other half
-
cell gives voltages which are the
standard potentials for those half
-
cells

E
cell

= 0 + E
half
-
cell

Zinc half
-
cell with SHE


Cell measures 0.76 V


Standard potential for Zn(s) = Zn
2+
(aq) + 2e = 0.76
V

Where there is no SHE


In this cell there is no SHE and the
measured voltage is 1.10 V

Standard reduction potentials


Any half reaction can be written in two ways:


Oxidation:

M = M
+

+ e (+V)


Reduction:

M
+

+ e = M (
-
V)


Listed potentials are standard

reduction
potentials

Applying standard reduction
potentials


Consider the reaction



What is the cell potential?


The half reactions are:



E
°

= 0.80 V


(
-
0.76 V) = 1.56 V


NOTE: Although there are 2 moles of Ag
reduced for each mole of Zn oxidized, we do not
multiply the potential by 2.

Extensive
v
intensive


Free energy is
extensive

property so need to
multiply by no of moles involved




But to convert to E we need to divide by no of
electrons involved




E is an
intensive

property


The Nernst equation


Working in nonstandard conditions

Electrode potentials and pH


For the cell reaction



The Nernst equation





Half
-
cell potential is proportional to pH

The pH meter is an electrochemical cell


Overall cell potential is proportional to pH






In practice, a hydrogen electrode is
impractical

Calomel reference electrodes


The potential of the calomel electrode is known vs
the SHE. This is used as the reference electrode
in the measurement of pH



The other electrode in a pH probe is a glass
electrode which has a Ag wire coated with AgCl
dipped in HCl(aq). A thin membrane separates
the HCl from the test solution

Cell potentials and equilibrium



Lest we forget…




So then



and

Cell potential a convenient way to
measure K

Many pathways to one ending


Measurement of K from different
experiments


Concentration data




Thermochemical data



Electrochemical data

Batteries


The most important application of galvanic
cells


Several factors influence the choice of
materials


Voltage


Weight


Capacity


Current density


Rechargeability

Running in reverse


Recharging a battery requires to run the
process in reverse by applying a voltage


In principle any reaction can be reversed


In practice it will depend upon many factors


Reversibility depends on kinetics and not
thermodynamics


Cell reactions that involve minimal structural
rearrangement will be the easiest to reverse

Lithium batteries


Lightweight (Molar mass Li = 6.94 g)


High voltage


Reversible process

Fuel cells


a battery with a
difference


Reactants are not contained within a sealed
container but are supplied from outside
sources

Store up not treasures on earth
where moth and rust…


An electrochemical mechanism for corrosion of iron. The metal and a
surface water droplet constitute a tiny galvanic cell in which iron is
oxidized to Fe
2+

in a region of the surface (anode region) remote from
atmospheric O
2
, and O
2

is reduced near the edge of the droplet at
another region of the surface (cathode region). Electrons flow from
anode to cathode through the metal, while ions flow through the water
droplet. Dissolved O
2

oxidizes Fe
2+

further to Fe
3+

before it is deposited
as rust (Fe
2
O
3
∙H2O).

Mechanisms


Why does salt enhance rusting?


Improves conductivity of electrolyte



Standard reduction potentials indicate which
metals will “rust”


Aluminium should corrode readily. It
doesn’t. Is thermodynamics wrong?


No, the Al
2
O
3

provides an impenetrable barrier

No greater gift than to give up your
life for your friend


A layer of zinc protects iron from oxidation, even when the
zinc layer becomes scratched. The zinc (anode), iron
(cathode), and water droplet (electrolyte) constitute a tiny
galvanic cell. Oxygen is reduced at the cathode, and zinc is
oxidized at the anode, thus protecting the iron from
oxidation.

Electrolysis


Electrolysis of a molten salt using inert electrodes


Signs of electrodes:


In electrolysis, anode is positive because electrons are removed
from it by the battery


In a galvanic cell, the anode is negative because is supplies
electrons to the external circuit

Electrolysis in aqueous solutions


a
choice of process


There are (potentially)
competing processes
in the electrolysis of an
aqueous solution


Cathode




Anode

Thermodynamics or kinetics?


On the basis of thermodynamics we choose
the processes which are favoured
energetically





But…chlorine is evolved at the anode

The role of
overpotentials


Thermodynamic quantities prevail only at
equilibrium


no current flowing


When current flows, kinetic considerations
come into play


Overpotential represents the additional
voltage that must be applied to drive the
process


In the NaCl(aq) solution the overpotential for
evolution of oxygen is greater than that for
chlorine, and so chlorine is evolved
preferentially


Overpotential will depend on the electrolyte
and electrode. By suitable choices,
overpotentials can be minimized but are never
eliminated


The limiting process in electrolysis is usually
diffusion of the ions in the electrolyte (but not
always)


Driving the cell at the least current will give
rise to the smallest overpotential

Electrolysis of water


In aqueous solutions of
most salts or acids or
bases the products will
be O
2

and H
2

Quantitative aspects of electrolysis


Quantitative analysis
uses the current
flowing as a measure
of the amount of
material


Charge = current x
time


Moles =
charge/Faraday