Mechanics

Oct 27, 2013 (4 years and 8 months ago)

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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

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

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”

doesn’t. Is thermodynamics wrong?

No, the Al
2
O
3

provides an impenetrable barrier

No greater gift than to give up your

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

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 =