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Electrochemistry
Engineering Chemistry
CHM 406
Electrochemical cell
Invented in early 1800’s (Volta)
 Daniell Cell shown (1830’s)

◦ Two half cells: Zn/Zn2+ and Cu/Cu2+
◦ “Salt bridge” connecting them internally
◦ External circuit
Current flows in the external circuit:
electrons
 Current is carried internally (aqueous
solution) by ions: Zn2+, Cu2+, K+, Cl-.

Half cell reactions
Anode: loss of electrons by zinc –
oxidation.
Zn (s) → Zn2+ (aq) + 2 e Cathode: gain of electrons by copper –
reduction.

Cu2+ (aq) + 2 e- → Cu (s)
 Overall reaction (redox reaction):
Zn (s) + Cu2+ (aq) → Zn2+ (aq) + Cu (s)
Electrolytes
When an electric current is passed through
an aqueous solution, the current is carried
by ions in the solution.
 Any substance which dissociates into ions
when dissolved in water is called an
electrolyte. A solution of an electrolyte
conducts electricity.
 Pure water itself is a very weak electrolyte –
contains traces of H3O+ and OH- ions.

Types of electrolytes

Salts: ionic compounds whose ions
dissolve in water.

Acids: covalent compounds which ionise
and dissociate when dissolved in water,
giving H3O+ (aqueous H+) ions.

Bases: molecules or ions which react with
acids, removing H+ ions.
Acids


Strong acids: completely dissociated.
Sulphuric acid: H2SO4(aq) → H+(aq) + HSO4-(aq)

Nitric acid: HNO3(aq) → H+(aq) + NO3-(aq)
Hydrochloric acid: HCl(aq) → H+(aq) + Cl-(aq)

Weak acids are partially dissociated.

Phosphoric acid: H3PO4(aq)
H+(aq) + H2PO42-(aq)
Bisulphate: HSO4-(aq)
H+(aq) + SO42-(aq)


Bases

React with acids, removing H+ ions, or react
directly with H+ (aq).

Hydroxide: OH-(aq) + H+(aq) → H2O (l)
Ammonia: NH3 (aq) + H+(aq) → NH4+ (aq)
Carbonate: CO32- (aq) + 2 H+(aq) → CO2(g) + H2O(l)



Compounds that dissociate into OH- ions in
water, e.g., NaOH, KOH, Ba(OH)2, are strong
bases.
Electrolysis
When an electric current is passed through
a solution of an electrolyte, or a molten salt,
electrolysis occurs.
 Different products are formed at each
electrode.
 Reverse of an electrochemical cell; electricity is used to force a non-spontaneous or
unfavoured reaction to occur.
 Can be used to produce substances
otherwise difficult to obtain, e.g., Na, Mg, Al.

Electrolysis of water
Usually done using a dilute solution of
H2SO4.
 Inert electrodes – platinum, graphite, etc.
 Cathode: reduction of H+.

◦ 2H+(aq) + 2e- → H2 (g)

Anode: H2O oxidised rather than SO42-.
◦ 2 H2O (l) → 4 H+(aq) + O2 (g) + 4e-

Overall:
◦ 2 H2O → 2 H2 + O2
Electrolysis of brine
Brine is concentrated aqueous NaCl, e.g.,
salt water.
 Cathode: reduction of H2O.

◦ 2 H2O (aq) + 2e- → H2 (g) + 2 OH- (aq)

Anode: Cl- oxidised rather than H2O.
◦ 2 Cl- (aq) → Cl2 (g) + 2e-

Overall:
◦ 2 H2O + 2 Cl- → H2 + Cl2 + 2 OH-

H2, Cl2, and aqueous NaOH are produced.
Electrolysis of molten salts and oxides
No water – no H2 formed.
 Used to manufacture metals which cannot
be obtained by chemical means.

◦ Highly reactive, e.g., Na (from NaCl), K (from
KCl), etc.
◦ Form very stable oxides, e.g., Al (from Al2O3), Ti
(from TiO2) etc.
Electroplating
Electrochemistry is used to deposit a layer
of a metal on an object, usually also made
of metal.
 The object to be plated must be de-greased
and thoroughly cleaned, then immersed in a
bath containing ions of the metal to be
plated, e.g., Cu2+, Ag+, Cr3+, etc.
 It is then connected to an external DC
source, such that it is the cathode.

Electroplating (contd.)
The anode is usually a piece of the metal to
be plated, though it can be inert.
 When the current is passed, the anode
gradually dissolves, while the cathode
(object to be plated) acquires a coating of
metal.
 E.g., silver (Ag) plating:

◦ Anode: Ag (s) → Ag+ (aq) + e◦ Cathode: Ag+ (aq) + e- → Ag (s)
Applications of electroplating
Aesthetic or decorative, e.g., gold or silver.
 Strengthening or protection against
corrosion, e.g., chromium, copper, zinc.
 Purification of impure metal: used as
anode, with a piece of pure metal as
cathode. The impure anode dissolves, and
pure metal is deposited on the cathode.

Redox reactions

Reactions in which electrons are
transferred from a reducing agent
(reductant) to an oxidising agent (oxidant).
◦ Combustion of fuels
◦ Corrosion of metals
◦ Conversion of metal ores to metals.
Examples of oxidants: O2, Cl2, metal cations,
metal oxides and oxy-anions (e.g., CrO42-).
 Examples of reductants: H2, compounds rich

in H and C, pure metals and cations with low
charge, anions such as Cl-, Br-, S2-, etc.
Half reactions

All redox reactions can be written as the
sum of two half reactions (whether or not
they are compartmentalised as in a cell).
◦ Oxidation half reaction.
◦ Reduction half reaction.

Example:
◦ Oxidation: Zn (s) → Zn2+ (aq) + 2 e◦ Reduction: Cu2+ (aq) + 2 e- → Cu (s)
◦ Overall: Zn + Cu2+ → Zn2+ + Cu
Common oxidation half-reactions






M (any metal) → Mn+ + n eH2 → 2 H+ + 2 eFe2+ → Fe3+ + e2 Cl- → Cl2 + 2 e2 S2O32- → S4O62- + 2 e2 H 2O → 4 H + + O2 + 4 e -
Common reduction half-reactions
Mn+ + n e- → M
 O2 + 4 e- + 4 H+ → 2 H2O
 F2 + 2 e- → 2 F Cu2+ + e- → Cu+


PbO2 + 4 H+ + SO42- + 2 e- → PbSO4 + 2 H2O
MnO4- + 8 H+ + 5 e- → Mn2+ + 4 H2O
 CrO42- + 8 H+ + 3 e- → Cr3+ + 4 H2O

Oxidation states (oxidation numbers)
How can we determine how many electrons
are gained or lost? Sometimes not obvious.
 The oxidation state of an atom is a number
that indicates the degree to which the atom
is oxidised (or reduced) compared to its
elemental state.
 The oxidation state of an elemental atom is
zero.
 The difference in oxidation states is the
number of electrons gained or lost.

Calculating oxidation states
The O.S. of a charged atom is equal to its charge.
 For covalently bonded atoms, polar covalent
bonds are treated like ionic bonds.
 F is always -1, except in F2.
 O is always -2, except when it is bonded to F or
another O.
 H bonded to C, N, O, or a halogen (Group 17) is
considered to be +1.
 The charge of a polyatomic ion is the sum of the
O.S.’s of its atoms.
 The O.S.’s of any other atom can be calculated
from the above.

Example: chromate
CrO42O.S. (Cr) + 4 x O.S.(O) = -2
O.S. (Cr) + 4 x (-2) = -2
O.S. (Cr) = +6
CrO42- (+6) → Cr3+ (+3) : 3 electrons needed.
Reaction takes place in acid; balance by
adding H+ to left hand side and H2O to right.
CrO42- + 8 H+ + 3 e- → Cr3+ + 4 H2O
Balancing redox equations
Multiply both oxidation and reduction
equations (must be balanced!) by numbers
such that the number of electrons gained
and lost is equal.
 If necessary, balance using H2O and either
H+ ions or OH- ions.
 Add the equations, so that the electrons
cancel out. Overall equation will be
balanced.

Example
MnO4- + Fe2+ → Mn2+ + Fe3+ (in acidic medium)
Oxidation: Fe2+ → Fe3+ + eReduction: MnO4- (+7) → Mn2+ (+2)
(5 electrons needed)
MnO4- + 8 H+ + 5 e- → Mn2+ + 4 H2O
5 Fe2+ → 5 Fe3+ + 5 eOverall: MnO4- + 8 H+ + 5 Fe2+ →
Mn2+ + 4 H2O + 5 Fe3+
Electromotive force / cell potential

For a current to flow in the external circuit
between two electrodes, there must be an
electromotive force. This emf is known as
the cell potential, E.
◦ Expressed in volts (V).
◦ Measured using a potentiometer or an ultra-low
current voltmeter.

The cell potential is the potential difference
between the two electrodes.
Electrode potentials
We can assign to every electrode (half cell)
a number in volts, called the electrode
potential, such that the difference between
two such numbers for two electrodes in a
cell, is the cell potential.
 The electrode potential is a measure of the
tendency of the half reaction at that
electrode (oxidation or reduction) to take
place.
 By convention, electrode potentials are
tabulated as reduction potentials.

Measuring electrode potentials
Absolute value cannot be measured.
 Values are measured relative to a standard
electrode.
 The Standard Hydrogen Electrode (S.H.E.)
is used as the reference electrode for this
purpose.
 By definition, the electrode potential for the
S.H.E. under standard conditions is zero.

Standard Hydrogen Electrode
Components of S.H.E.






Inert electrode (Pt foil – H2 easily adsorbed
on the surface).
Immersed in dilute H2SO4, such that H+
concentration is 1 mol L-1.
H2 gas bubbled into the solution past the
electrode, such that the H2 pressure is 1 bar
(105 Pa).
Temperature is 25oC (298 K)
Salt bridge connects the half cell to the other
half cell, whose potential is to be determined.
2 H+ (aq) + 2 eH2 (g) ; Eo = 0 V
Standard Conditions
Temperature = 25oC (298 K)
 Partial pressure of every gaseous reactant
or product = 1 bar (105 Pa)
 Concentration of every aqueous reactant or
product = 1 mol L-1
 Reduction potentials measured under these
conditions, relative to the S.H.E. are called
“standard reduction potentials,” Eo.

Standard Reduction Potentials
Li+ (aq) + eLi (s)
Mg2+ (aq) + 2eMg (s)
Al3+ (aq) + 3eAl (s)
Zn2+ (aq) + 2eZn (s)
Cr3+ (aq) + 3eCr (s)
Fe2+ (aq) + 2eFe (s)
Ni2+ (aq) + 2eNi (s)
2 H+ (aq) + 2eH2 (g)
Cu2+ (aq) + 2eCu (s)
I2 (s) + 2e2 I- (aq)
Ag+ (aq) + eAg (s)
O2 (g) + 4H+ + 4e2 H2O (l)
F2 (s) + 2e2 F- (aq)
Eo
Eo
Eo
Eo
Eo
Eo
Eo
Eo
Eo
Eo
Eo
Eo
Eo
=
=
=
=
=
=
=
=
=
=
=
=
=
- 3.04 V
- 2.38 V
- 1.66 V
- 0.76 V
- 0.74 V
- 0.41 V
- 0.23 V
0.00 V
0.34 V
0.54 V
0.80 V
1.23 V
2.87 V
Electrochemical series





Lists half reactions in order of reactivity.
More positive Eo  greater tendency of
the reaction to occur.
For oxidation potentials – reverse the
reaction & change the sign of Eo.
Half reactions at the top of the series are
more likely to occur in reverse: best reducing
agent is Li.
Half reactions at the bottom are more likely
to occur as shown: best oxidising agent is F2.
Types of electrodes

Reactive metals in contact with solutions of
their cations.
◦ E.g., Zn(s) | aq. ZnSO4

Reactive gases bubbling through solutions of
cations or anions, in contact with an inert
electrode.
◦ E.g., Pt(s) | Cl2(g) | aq. KCl

Inert electrodes (Pt) in contact with aqueous
solutions of the reactants & products of the half
reaction
◦ E.g., Pt(s) | Fe2+(aq), Fe3+(aq)
Types of electrodes (contd.)

Reactive metals coated with an insoluble
salt of that metal, the salt in contact with a
solution of its anions.
◦ E.g., Ag(s) | AgCl(s) | aq. KCl
AgCl (s) + eAg (s) + Cl- (aq)

Membrane electrodes.
 E.g., glass electrode, used to measure
concentration of H+ ions in solution (pH).
The calomel electrode
The calomel electrode (contd.)
The S.H.E. is not a convenient electrode for
regular use as a reference.
 A reference electrode needs to be easy to
use, stable, reproducible, and reliable.
 The calomel electrode is the electrode of
choice: Pt (s) | Hg (l) | Hg2Cl2 (s) | saturated
aq. KCl.

Hg2Cl2(s) + 2e2 Hg(l) + 2 Cl-(aq)
 E = + 0.2444 V (relative to S.H.E.)

Standard Cell Potentials
In principle, any two half-cells can be
combined to give a cell.
 The cell potential Ecell is the difference
between two reduction potentials, or
equivalently, the sum of a reduction and an
oxidation potential.
 If determined under standard conditions, it
is a standard cell potential, Ecello.
 Ecell positive  spontaneous reaction.

Electrochemical work
When an electric current is passed between
2 electrodes, the electrical work done is
w = - q x Ecell
 q = charge (coulombs; q = i (in A) x t (in s) ).

For a redox reaction, q can be determined
in terms of moles of electrons transferred.
q=nF
 n = number of moles of electrons
 F = charge per mole of electrons

Electrochemical work (contd.)
F is known as Faraday’s constant.
 F = 96,480 C mol-1 (J V-1 mol-1).
 The work that can be done is referred to as
free energy, DG.
 Therefore
DG = - n F Ecell

The Nernst Equation

The Nernst equation is used to determine
Ecell when conditions are not standard, i.e.,
concentrations ≠ 1 mol L-1.
Ecell  E

0
cell
2.303 R T

log Q
nF
Q = quotient
 R = gas constant = 8.314 J K-1 mol-1
  F = Faraday’s constant
 n = number of electrons transferred.
What is the “quotient?”
Consider any reaction (redox or otherwise), which may
be denoted
aA + bB + cC + …
xX + yY + zZ + …
where A, B, C, etc. are the reactants / products,
and a, b, c, etc. are the coefficients in the balanced
equation.
Let [A], [B], [C], etc., be the concentrations of A, B, C,
etc., in mol L-1.
Then
[X] x [Y] y [Z] z ...
Q =
[A]a [B] b [C] c ...
Thus, if the various concentrations are known, Ecell can
be calculated.
