Electrochemical cells 1 Overview
•  Redox reac2ons •  Electrochemical cells •  Electrode poten2al and cell voltage (electromo2ve force) •  Voltaic cells •  Standard poten2als •  Non-­‐standard poten2als (first look) 2 Redox reactions
•  Oxida2on –  loss of electrons H2 ! 2H+ + 2e
•  Reduc2on –  gain of electrons O2 + 4H+ + 4e ! 2H2 O
•  In an electrochemical cell, oxida2on occurs on an electrode called an anode in a half cell •  Reduc2on occurs on an electrode called an cathode in a second half cell •  Reduc2on occurs spontaneously on the electrode with the highest poten2al (called the posi2ve electrode). •  Oxida2on happens spontaneously on the nega2ve electrode (lowest poten2al) •  The number of electrons transferred (n) must balance in each half cell •  The general formula for a ‘simple’ oxida2on-­‐reduc2on reac2on is νOO+ne- ⇔ νRR
•  O is the oxidized state (oxidizing agent) and R is the reduced state (reducing agent). n is the number of electrons transferred 3 Redox Reactions
•  Direct redox reac2on –  Oxidizing and reducing agents are mixed together (everything happens on one electrode, which has a certain electrical poten2al) •  Indirect redox reac2on –  Oxidizing and reducing agents are separated (in half cells) but connected electrically –  Basis for electrochemical cell –  E.g. H2 + (1/2)O2àH2O, which is an overall reac2on made up of two charge transfer reac2ons occurring in two half cells on electrodes –  H2 is the reducing agent and O2 is the oxidizing agent 4 Electrode potential
zinc electrode
Zn2+ ions in solution
!
Zn(s) ↵ Zn2+ (aq) + 2e
•  zinc rod (electrode) immersed in a solu2on (electrolyte) containing Zn2+ ions •  Reac2on occurs both ways un2l charge build up in the solu2on prevents further change •  A poten2al difference is set up between the rod and the solu2on: called the electrode poten2al A simple electrochemical cell
•  Zinc rod (electrode) immersed Solu2on becomes in a solu2on containing Zn2+ -ve
+ve
V
less + ve/more ions -­‐ve so need to SO42-­‐ transfer nega2ve •  Connect to a copper rod charge immersed in a solu2on Zn
Cu
containing Cu2+ ions •  A poten2al difference is set up between the electrodes (difference in their poten2als) •  Put a salt bridge between the Cu2+
Zn2+
solu2ons !
•  Nega2vely charged sulfate ions Zn(s) à Zn2+(aq) + 2e Cu2+(aq) + 2e à Cu(s) (SO42-­‐) pass through the salt (oxida2on: anode) (reduc2on: cathode) bridge to main electroneutrality Overall: Zn(s) + Cu2+(aq) à Zn2+(aq) + Cu(s) in both half cells 6 Electrochemical Cells (2 basic types)
•  Voltaic Cell –  cell in which a spontaneous redox reac2on generates electricity. Electrons flow from a nega2ve to a posi2ve terminal, i.e., from anode to cathode since they carry a nega2ve charge and will be repelled by nega2ve charge and a]racted by posi2ve charges –  Note that conven2onal current is defined in the opposite way –  Electrons flow spontaneously from low to high poten2al (which is defined for posi2ve charges) so they go from high to low poten2al energy –  Spontaneity also in terms of energy of reac2on (later) –  chemical energy to converted to electrical energy –  Fuel cells and ba]eries (when discharging) are voltaic cells •  Electroly2c Cell (electrolysis) –  electrochemical cell in which an electric current drives a nonspontaneous redox reac2on –  Force electrons to the nega2ve (lower poten2al) electrode and draw electrons from the higher poten2al electrode –  Electrolysis is used to drive an oxida2on-­‐reduc2on reac2on in a direc2on in which it does not occur spontaneously –  E.g. electrolysis of water 2H2O à 2H2 + O2 7 EMF
•  Electrically connec2ng 2 electrodes (in half cells) leads to an electrical poten2al difference between the electrodes, which may allow a current to flow between them •  In a voltaic cell electrons are produced on the –ve electrode and are driven to the +ve to lower their poten2al energy •  In an electroly2c cell we have to drive the reac2ons and therefore the flow of electrons from the posi2ve to nega2ve electrode (low to high poten2al energy) •  The poten2al difference from the cathode to anode is called an electromo2ve force (EMF) or cell voltage (wri]en E or
Ecell) •  EMF is driving force for electron (i.e. a current) flow between the two electrodes from high to low poten2al •  Must allow ions to flow from half cell to half cell (by a salt bridge, separator or membrane) to maintain charge neutrality 8 Voltaic Cells
•  Voltaic (Galvanic) produce electrical power by separa7ng an oxida7on and reduc7on process in half cells that are electrically connected for current flow and allow internal transfer of ions to maintain charge balance (electroneutrality) •  For the process to be spontaneous the poten7al of the electrode where reduc7on takes place must be higher that the oxida7on poten7al 9 Voltaic Cells: General
Nega2ve electrode (anode) Electron flow electrolyte Oxida2on Posi2ve electrode (cathode) Reduc2on Nega2ve half Posi2ve half cell cell (anode) separator 10 Voltaic Cells: General
•  The overall reac2on is achieved by two reac2ons, one in each half cell o See previous example o H2/O2 fuel cell: Overall: H2 + (1/2)O2àH2O −ve electrode (anode) H2 à 2H+ + 2e− +ve electrode (cathode) (1/2)O2 + 2H+ + 2e− à H2O •  Each half cell reac2on is a reduc2on-­‐oxida2on (redox) reac2on occurring predominantly in one direc2on (except when no current flows) 11 Hydrogen/Oxygen fuel cell
•  Hydrogen is oxidised in the anode •  Oxygen is reduced in the cathode •  Only product is water Overall reac2on: H2 + (1/2)O2àH2O Anode: H2à2H++2e− Cathode: (1/2)O2+2H++2e−àH2O 12 Primary ba]eries •  Ba]eries such as dry cells, Zinc air alkaline cells and bu]on ba]ery cells are non-­‐rechargeable. •  Cells that cannot be recharged are called primary cells or ba]eries •  Example is the zinc-­‐air Anode: ba]ery shown on the lef Zn + 2OH-­‐ -­‐ 2e-­‐ è Zn(OH)2 Cathode: •  Typically small and used in hearing aids and watches 1/2 O2 + H2O + 2e-­‐ è 2OH-­‐ •  Many others, alkaline, zinc-­‐
Overall reac2on: carbon, Al-­‐air 2Zn + O → 2ZnO 2
13 The Lead Acid (secondary) Ba]ery •  This is a rechargeable or secondary ba]ery •  Two electrodes, one of lead, the other of lead dioxide (PbO2) immersed in sulfuric acid •  Lead is oxidised and lead ions (Pb2+) dissolve, leaving two electrons behind at the anode •  Two electrons flow through the circuit and are used to reduce lead dioxide •  Flow of protons from anode to cathode to maintain charge balance 14 Standard electrode potentials
•  The electrode poten2al depends on the local condi2ons (including temperature, pressure and concentra2on of reactants) •  Electrode poten2als measured under standard condi2ons (usually 298 K, 1 atm and with all species in solu2on 1.0 mol dm−3) are called standard electrode poten2als, denoted E0. •  Eg., Fe2+(aq) + 2e− = Fe(s) has E0 = −0.44 V •  This value will change with condi2ons (later) •  It is important to note that these poten2als apply only to Nernst equilibrium condi2ons (no net current) •  When there is a net transfer, an overpoten2al is established (more on this later) 15 Standard electrode potentials
•  The EMF is easy to measure, but the individual electrode poten2als themselves cannot actually be measured at all (it is only possible to measure the poten2al difference between two electrodes) •  Solu2on: Arbitrarily choose a value for some electrode (called a reference electrode). Measure all other half cell poten2als rela2ve to this value •  In detail, put the electrode of interest in a cell with the reference electrode (where oxida2on is forced) and measure the EMF •  Not really an issue because we are only interested in the poten2al difference between electrodes •  The electrode conven2onally used for this purpose is the standard hydrogen electrode H
2 ↵
2H
+ +
2e
, arbitrarily assigned an E0 value of 0.0V 16 Standard Reduc2on Poten2als The half cell poten2als E0 that are determined by reference to the SHE (trea2ng it as the anode) More oxidizing More reducing 17 Standard Reduc2on Poten2als and Ecell •  The more +ve the standard poten2al, the more oxidizing the species •  The more -­‐ve the standard poten2al, the more reducing the species •  To find the cell voltage or EMF under standard condi7ons E0cell = E0cathode – E0anode or E0cell = E0reduction – E0oxidation •  This is only valid under standard condi2ons (hence ‘0’) and is a theore2cal ideal value (maximum a]ainable) that is a]ained when there is no net transfer of charge, called the open circuit voltage (OCV) •  For a Galvanic cell, the half cell reac2on with the more posi2ve E0 will be the cathode (because this ac2ve species is more oxidising so will oxidise the species in the other half cell (which is therefore an anode) •  A posi2ve E0cell means reac2ons occur spontaneously, i.e., it is a voltaic cell (if <0 then it is an electroly2c cell) – the reason for this is thermodynamic, which will become clear when we look at Gibb’s free energy change. It is also (as we have seen) a consequence of electrons lowering their P.E. 18 Example 1: What is the standard EMF of a galvanic cell made of half cells with the following reac2ons under standard condi2ons Cd2+(aq) + 2e− ↵
Cd(s) E0 = −0.40 V Cr3+(aq) + 3e− ↵
Cr(s) E0 = −0.74 V •  Cd poten2al is more posi2ve, therefore it will oxidize Cr, therefore it will occur in the cathode (reduc2on) •  Note that the reac2ons will be (to balance the electrons) 3Cd2+(aq) + 6e− ↵
3Cd(s) (cathode) 2Cr(s) ↵
2Cr3+(aq) + 6e− (anode) E0cell = E0cathode – E0anode = – 0.4 – (–0.74) = 0.3 V
0 = _________
Ecell
19 Example 2: What is the standard EMF of an electrochemical cell made of half cells with the following reac2ons under standard condi2ons Cu2+(aq) + 2e− ↵
Cu(s) E0 = +0.34 V Zn2+(aq) + 2e− ↵
Zn(s) E0 = −0.76 V •  Cu poten2al is more posi2ve, therefore it will oxidize Zn, therefore it will occur in the cathode (reduc2on) •  Note that the reac2ons will be (to balance the electrons) Cu2+(aq) + 2e− ↵
Cu(s) (cathode) Zn(s) ↵
Zn2+(aq) + 2e− (anode) E0cell = E0cathode – E0anode = 0.34 – (–0.76) = 1.1 V
0 = _________
Ecell
20 Other condi2ons •  For non-­‐standard condi2ons, electrode poten2als and therefore the (open-­‐circuit) voltage will change •  The OCV is ofen called the reversible OCV (ROCV) or thermodynamically derived ROCV •  When a current flows through the cell it leads to irreversible losses due to limita2ons in mass transport, charge transport and energy barriers associated with the charge transfer reac2ons •  To study these properly we will need to look at the thermodynamics and electrochemical kine2cs of voltaic cells •  The cell voltage in that case will be the ROCV (under the relevant condi2ons) minus several losses 0 = _________
Ecell
21 able 10.1,
the half-reactions
of
common
oxidants and
reductants
are
listed.
Table
10.1
Somemagnesium,
examples
ofisoxidants
oxidants
and
reductants
hlorine,
which
gained
electrons
from
a accept
reductant.
Generally,
suc
the counterpart,
it can
be a10.1
reductant.
If it can easily
electron(s),
it can
Table
Some
examples
of
and
reductants
xidants
ements
alkali
metals orofalkaline
earth
metals
strong are
reductants;
able
10.1,asthe
half-reactions
common
oxidants
and are
reductants
listed. such
xidants
Quiz
-- Table --10.1 Some
of oxidants and reductants
(aq) +
+ 2e
2eare good
2I (aq)
(aq)
ements asIIchlorine
oxidants. examples
22(aq)
2I
xidants
-- r
--(aq)
A1.
compound
can
act
as
an
oxidant
and
as
a reductant
as swell.
Ifaitnd can easily d
Iden2fy t
he educed state, the oofxidised tate Table
10.1
Some
examples
oxidants
and reductants
Br
(aq)
+
2e
2Br
2
Br
(aq)
+
2e
2Br
(aq)
2
I2(aq) 2-(
+ it2ecan be2Ia++(aq)
2-(
3+ easily accept electron(s), it can
the counterpart,
reductant.
it2Cr
can
xidants
-- If
3+
O
aq)
+
14H
(aq)
+
6e
(aq) ++ 7H
7H
Cr
the n
umber o
f e
lectrons t
ransferred in 22O(l)
tO(l)
he 2
7
- 14H -(aq)
- + 6e
O
aq)
+
2Cr
(aq)
Cr
2
7
+2e2e-- 2I2Br
2(aq)
- (aq)
able 10.1,Br
the
half-reactions
of
common
oxidants and reductants are listed.
ICl
+
(aq)
2(aq)
(aq)
+
2e
2Cl
(aq)
2
2-(
+
3+
following: Cl2(aq)
+
2e+- 14H2Cl
(aq)
O
aq)
(aq)
+
6e
2Cr
Cr
22(aq)
7 - + 2e
+ 2Br (aq) -2+ (aq) + 7H2O(l)
Br
+
2+
MnO
(aq) +
+- 8H
8H (aq)
(aq)
+ 5e
5e - Mn
Mn (aq)
(aq)
+ 4H
4H22O(l)
O(l)
-+
MnO
(aq)
442-(
+10.1
3+ of+oxidants
Table
Some
examples
and reductants
Cl
(aq)
+
2e
2Cl
(aq)
22O222aq)
+
14H
(aq)
+
6e
2Cr
(aq)
+
7H
O(l)
Cr
7
2
2O88 -(aq)
(aq) +
+- 2e
2e + 2SO
2SO
(aq) 2+
S
S22O
xidants MnO
44 - (aq)
+ 8H (aq)
+
5e
Mn (aq) + 4H2O(l)
4 (aq)
Cl2(aq)
+ 2e
2Cl
(aq)
eductants
eductantsI2(aq)2--+ 2e
2- 2I
+ (aq)
- (aq)
2+
2+
O
(aq)
+
2e
2SO
S
2
8
4
MnO
(aq)
+
8H
(aq)
+
5e
Mn
(aq) + 4H2O(l)
2+
4
Zn(s)
Zn
(aq)
+
2e
- (aq)
-2e
Zn(s)
Zn
+
2- + 2e + - 2Br (aq)
22(aq)
eductants Br
O
(aq)
+
2e
2SO
S
+
22(g)
8 2-(
4 (aq)
H
2H
(aq)
+
2e
++
3+
H
(g)
2H
(aq)
2e
2+
2
O
aq)
+
14H
(aq)
+
6e
2Cr
(aq) + 7H2O(l)
Cr
2
7
Zn(s)
Zn
(aq)
+
2e
+
eductants H S(aq)
+ (aq) -+ S(s) + 2e2H
- 2H
2S(aq)
H
(aq)
+(aq)
+ 2e
+
-S(s)
2
2+
Cl
(aq)
+
2e
2Cl
2
H
(g)
2H
(aq)
+
2e
2
Zn(s)
Zn (aq)
2+
4+ + 2e
4+
-Sn2+
(aq)
Sn
(aq)
+
2e
+
- 2+(aq) + 4H O(l)
+ + (aq) + -5e
MnO
(aq)
+
8H
Mn
4
2
H22+
S(aq)
2H
S(s)
+
2e
(g)
2H (aq)
2+
3+ + 2e -3+
Fe 2+(aq)
Fe- 4+
(aq) + e 2-2+
O
(aq)
+
2e
S
2
8
4 (aq)
Sn
(aq)
Sn (aq)
(aq)2SO
2e
H2S(aq)
2H
++S(s)
+ 2e2+
3+
- 2+
4+
eductants Fe
(aq)
Fe
(aq)
+
e
Sn
Sn
(aq)
+
2e
number2+
e)) Oxidation
2+
3+ + 2e Zn(s)
Zn
(aq)
2+
3+
2+
3+
(aq)reductant
Fe (aq)
+
e
Iron isFe
a good
and
it
becomes
Fe
or
Fe
depending on
on the
the reaction
reaction
depending
0 -= _________
+
E
H2(g)number
2H (aq) +cell
2e
) Oxidation
+ and it becomes- Fe2+ or Fe3+ depending on the reaction
Iron
is
a
good
reductant
2+
22 H2S(aq)
2H (aq) + S(s) + 2e
) Oxidation
number
Fe
Fe +2e
2+
3+
are -3.045 V and +2.87 V,Quiz
respectively.
You can understand th
2. electromotive
In the previous examples, state whether is a and fluo
force.
A combination
ofthe Li reac2on electrode
reduc2on or oshould
xida2on whether it occurs safety
in an aof
node or cells
Evidently
care
beand taken
to ascertain
such
cathode alkali
metal/alkali metal electrode is in fact used for the alkal
3. widely
Write used.
the overall reac2ons for the cells in Examples 1 and 2 4.  Find the cell voltage or EMF under standard condi2ons in Sample
exercise
10.5 Electromotive
forces
ofthe cells
cells with the following overall reac2ons (use list of Calculate
the standard
force of cells (25 °C)
reduc2on poten2als on the electromotive
next slide) 10.2.
Sn2+ + Pb
(1) Sn + Pb2+
(2) 2Fe3+ + Sn2+
2Fe2+ + Sn4+
(3) 5Fe2+ + MnO4- + 8H+
5Fe3+ + Mn2+ + 4H2O
Answer
(1) 0.009 V
(2) 0.617 V (3)
0.739 V
0 = _________
Ecell
(c) Nernst’s equation
23 The concentration dependence of electrode potentials h
I + 2e- 2I0.535
The normal electrode potentials of important electrodes are given in2 Table
+
-10.2.
Cu + e
Cu
0.521
O2 + 2H225°C)
O + 4e
4OH
0.401
Table 10.2 The normal electrode potential V (as aqueous solution,
2+
Cu + 2e
Cu
0.337
AgCl + e
Ag + Cl
0.222
F2 + 2e
2F
2.87
2+
3+
2+
SO4 + 4H + 2e
H2SO3 + H2O
0.171
Co
+e
Co
1.92
4+
2+
+
Sn + 2e
Sn
0.154
H2O2 + 2H + 2e
2H2O
1.776
2+
+
+
Cu + eCu
0.153
MnO4 + 4H + 3e
MnO2 + 2H2O
1.695
+
+
22H + 2e
H2
0.00
PbO2 + 4H + SO4 + 2e
PbSO4 + 2H2O
1.685
2+
+
2+
Pb + 2e
Pb
-0.129
MnO4 + 8H + 5e
Mn + 4H2O
1.51
2+
3+
Sn + 2e
Sn
-0.138
Au + 3e
Au
1.50
2+
+
2+
Ni + 2e
Ni
-0.228
PbO2 + 4H + 2e
Pb + 2H2O
1.455
2Pb + SO4
-0.355
PbSO4 + 2e
2Cl
1.396
Cl2 (aq) + 2e
2+
2+
3+
Cd + 2e
Cd
-0.402
Cr2O7 + 14H + 6e
2Cr + 7H2O
1.29
3+
2+
+
Cr + e
Cr
-0.424
O2 + 4H + 4e
2H2O
1.229
Ch 10 Oxidation and reduction
2+
+
2+
Fe + 2e
Fe
-0.440
MnO2 + 4H + 2e
Mn + 2H2O
1.23
Br2 (aq) + 2e
2Br
1.087
Cr3+ + 3e- Cr
-0.67
3+
2+
NO + 4H + 3e
NO + 2H2O
0.957
Zn + 2e
Zn
-0.763
2+
2+
10(19)
2Hg + 2e
Hg2
0.920
H2(g) + 2OH
-0.828
2H2O + 2e
2+
Ag+ + e- Ag
0.799
Mn + 2e
Mn
-1.18
2+
3+
Hg2 + 2e
2Hg
0.789
Al + 3e
Al
-1.662
3+
2+
Fe + e
Fe
0.771
2H
-2.25
H2 + 2e
+
2+
O2 + 2H + 2e
H2O2
0.682
Mg + 2e
Mg
-2.37
2+
MnO4 + e
MnO4
0.558
Na + e
Na
-2.714
2+
I2 + 2e
2I
0.535
Ca + 2e
Ca
-2.84
+
2+
Cu + e
Cu
0.521
Ba + 2e
Ba
-2.92
+
O2 + 2H2O + 4e
4OH
0.401
K +e
K
-2.925
2+
+
Cu + 2e
Cu
0.337
Li + e
Li
-3.045
AgCl + e
Ag + Cl
0.222
2+
Judging from the examples above, it is expected that electrodes m
SO4 + 4H + 2e
H2SO3 + H2O
0.171
0
4+
2+
cell
Sn + 2e
Sn
0.154 ionization tendency will have a large, negative normal electrode potent
2+
+
Cu + eCu
0.153 halogens with large electronegativity will have a large, positive normal e
indeed the case, and the normal electrode potentials of these electrode
rea
24 2H+ + 2e- H2
0.00
Quiz E
2+
-
= _________
+