This document provides an overview of electrochemistry and some key concepts. It begins by defining electrochemistry as the study of how spontaneous chemical reactions can produce electricity and how electrical energy can drive non-spontaneous reactions. It then discusses several applications of electrochemistry including metal production, electroplating, and batteries. The document goes on to define conductors and the differences between metallic and electrolytic conduction. It also introduces concepts like galvanic cells, salt bridges, standard electrode potentials, and the electrochemical series. In summary, the document provides a broad introduction to fundamental electrochemistry topics and concepts.

1. ELECTROCHEMISTRY
Class XII
By: ARUNESH GUPTA
PGT (CHEMISTRY)
KendriyaVidyalaya Barrackpore (AFS)

2. Contents:
Introduction to electrochemistry & its applications
Metallic & electrolytic conduction.
Types of electrolytes
Electrochemical cell & electrolytic cell
Galvanic cell ( Daniell cell )
Salt bridge
Electrode potential: reduction potential & oxidation potential.
Standard states, standard reduction potential, standard oxidation potential
Reference electrode (SHE, Calomel electrode)
Standard e.m.f. & standard cell potential & their difference
Electrochemical series & its applications
Nernst equation
Standard cell potential & equilibrium constant
Standard cell potential & standard Gibb’s energy
Numerical problems

3. Unit – 3 Electrochemistry
• Electrochemistry is that branch of chemistry which deals with the
study of production of electricity from energy released during
spontaneous chemical reactions and the use of electrical energy to
bring about non-spontaneous chemical transformations.
• Importance of Electrochemistry
• (1) Production of metals like Na, Mg, Ca and Al etc.
• (2) Electroplating.
• (3) Purification of metals.
• (4) Batteries and cells used in various instruments.
• (5) Explanation of natural phenomenon like corrosion etc.

4. Conductors : Substances that allow electric current to pass
through them are known as conductors.
(i) Metallic Conductors or Electronic Conductors: Substances which allow the electric current to
pass through them by the movement of electrons are called metallic conductors, e.g.. metals.
It depends on
(i) the nature & structure of the metal
(ii) the no. of valence electrons per atom
(iii) the density of metal & (iv) temperature (it decreases with the increase in temperature)
(ii) Electrolytic Conductors or Electrolytes: Substances which allow the passage of electricity
through their fused state or aqueous solution and undergo chemical decomposition are called
electrolytic conductors, e.g., aqueous solution of acids. bases and salts.

5. Difference between metallic & electrolytic conductions
Metallic Conduction Electrolytic conduction
1) Metallic conduction is carried by the
movement of electrons
1) Electrolytic conduction is carried by the
movement of ions.
2) It involves no change in the chemical
properties of the conductor.
2) It involves the decomposition of the
electrolyte as a result of the chemical
reaction.
3) It does not involve the transfer of
matter or ions.
3) It involves the transfer of matter or ions.
4) Metallic conduction decreases with
the increase in temperature.
4) Electrolytic conduction increases with
the increase in temperature.

6. Types of Electrolytes:
Electrolytes are of two types:
1. Strong electrolytes: The electrolytes that completely dissociate
or ionise into ions are called strong electrolytes. e.g., HCl,
NaOH, K2SO4 etc.
2. Weak electrolytes
The electrolytes that dissociate partially having degree of
dissociation less than 1 are called weak electrolytes, e.g.,
CH3COOH, H2CO3, NH4OH, H2S, etc.

8. An electrochemical is the arrangement in which chemical energy of a
spontaneous reaction is converted into electrical energy.
An electrochemical cell of almost constant emf is called standard cell.The
most common is Weston standard cell. Galvanic cell is also called voltaic cell.
General Representation of an Electrochemical Cell :
M1(s) / M1
n+ (aq) || M2
n+(aq)/M2(s)
Anode→ 𝑶𝒙𝒅𝒏. Redn.← Cathode
Other features of the electrochemical cell are
(1)There is no evolution of heat.
(2)The solution remains neutral on both sides.
(3)The reaction and flow of electrons stops after sometime.
The electrochemical cell

9. (a) When Ecell <1.1 V (b) When Ecell = 1.1 V (c) When Ecell > 1.1 V
(i) Electron flow from Zn
rod to Cu rod hence
current flows from Cu to
Zn.
(i) No flow of electrons or
current
(i) Electrons flow & current
flows from Zn to copper.
(ii) Zn dissolves at anode
and Cu deposits at
cathode. Cell reaction is
takes place.
(ii) No chemical reaction. (ii) Zn is deposited at the
Zn electrode & Cu is
dissolved at Cu electrode.
Cell reaction is reversed.
When a Daniell cell is externally connected with a cell of different cell potential Ecell

10. Electrochemical Cell and
Electrolytic Cell
A spontaneous chemical process is the one which can take place on its
own and in such a process the Gibb’s energy of the system decreases. It is
this energy that gets converted to electrical energy.
The electrochemical cell converts chemical energy of a spontaneous
reaction into electrical energy.
The reverse process is also possible in which we can make non-
spontaneous processes occur by supplying external energy in the form of
electrical energy. These inter conversions are carried out in equipment
called Electrolytic Cells.
The reactions carried out electrochemically can be energy efficient and
less polluting.

11. GALVANIC CELL (Daniell Cell)
 Daniell Cell is a Galvanic Cell in which Zinc
and Copper are used for the redox reaction to
take place.
 Zn is the reducing agent and Cu2+ is the
oxidising agent. The half cells are also known as
Electrodes.
 The oxidation half is known as Anode and the
reduction half is called Cathode.
 Electrons flow from anode to cathode in the
external circuit. Anode is assigned negative
polarity and cathode is assigned positive
polarity.

12. Daniell cell
 In Daniell Cell, Zn acts as the anode and Cu acts as the cathode.
 At Anode (oxidation) Zn(s)  Zn2+(aq) + 2e
 At Cathode (reduction) Cu2+(aq) + 2e  Cu(s)
 Cell reaction: Zn(s) + Cu2+(aq)  Zn2+(aq) + Cu(s)
 The cell is represented as: Zn(s) / Zn2+(aq) // Cu2+(aq) / Cu(s)
 AID to remember  Representation of a cell
 A-B-C (A Anode, B Bridge, C Cathode
 O//R (O Oxidation reaction // R Reduction reaction
 In general : M1(s)/M1
a+
C1 || M2
b+
(C2)/ M2(s).
 Standard cell potential Ecell = 1.10 V

13. Electrochemical cell Anode Cathode
Position (L) Left electrode Right electrode
Reaction (O) Oxidation Reduction
Name of electrode (A) Anode Cathode
Charge of electrode (N) Negative electrode Positive electrode
Movement of electrons Out of cell from anode Into the cell from cathode
By convention an anode on the LHS and cathode is represented on the RHS. [REMEMBER: L O A N ]

14. Salt bridge
• Salt bridge: A salt bridge connects the two electrolytic solutions of an electrochemical cell.
• Functions of a salt bridge:
• 1) It completes the circuit and allows the flow of current.
2) It maintains the electrical neutrality on both sides. Salt-bridge generally contains solution of strong
electrolyte such as KNO3, KCl etc in agar-agar or gelatine gel etc.
• (Jelly does not allow the salt of salt bridge to mix with the electrolytic solution).
• KCI is preferred because the transport numbers of K+ and Cl-are almost same. (Ionic mobility of K+(aq) & Cl-
(aq) is nearly same so it reduces the internal resistance of a cell.
Preferential Discharge of ions :
• Where there are more than one cation or anion the process of discharge becomes competitive in nature.
Discharge of any ion requires energy and in case of several ions being present the discharge of that ion will
take place first which requires the energy.

15. Electrode Potential (𝐸𝑒𝑙𝑒𝑐𝑡𝑟𝑜𝑑𝑒)
• When an electrode is in contact with the solution of its ions in a half-
cell, it has a tendency to lose or gain electrons which is known as
electrode potential. It is expressed in volts.
• It is an intensive property, i.e., independent of the amount of species
in the reaction.
• So, the potential difference between an electrode and its ion in
solution is called its electrode potential (𝐸𝑒𝑙𝑒𝑐𝑡𝑟𝑜𝑑𝑒)
• The electrode potential will be named as oxidation or reduction
potential depending upon whether oxidation or reduction has taken
place.

16. Oxidation potential (Eoxdn or 𝑬 𝑴(𝒔)/𝑴 𝒏+)
The tendency to lose electrons when an electrode is in contact with its ions is known as oxidation potential. Oxidation
potential of a half-cell is inversely proportional to the concentration of ions in the solution. It is denoted by 𝐸 𝑀/𝑀 𝑛+
The half reaction is M ⇌ Mn+ + ne (Oxidation)
Reduction potential (Eredn or 𝑬 𝑴 𝒏+/𝑴(𝒔))
The tendency to gain electrons when an electrode is in contact with its ions in the solution is known as reduction
potential.
It is denoted by 𝐸 𝑀 𝑛+/𝑀 . The half reaction is Mn+ + ne ⇌ M (reduction)
According to IUPAC convention, the reduction potential alone be called as the electrode potential unless it is
specifically mentioned
Remember: (a) Both oxidation and reduction potentials are equal in magnitude but opposite in sign.
(b) It is not a thermodynamic property, so values of Eredn are not additive.
Ereduction = – Eoxidation
It is not possible to determine the absolute value of electrode potential, as oxidation & reduction takes place
simultaneously & not individually. So, a reference electrode [NHE or SHE] is required to determine its value.
The electrode potential is only the difference of potentials between two electrodes that we can measure by
combining them to form a complete cell.

17. Standard Electrode Potential (𝑬 𝒐
𝒐𝒓 𝑬 𝜽
)
The potential difference developed between metal electrode and solution of its ions of unit molarity (1M) at 1
atm pressure and 25°C (298 K) is called standard electrode potential. It is denotes as 𝐸𝑒𝑙𝑒𝑐𝑡𝑟𝑜𝑑𝑒
𝑜
. It can be
standard reduction potential or standard oxidation potential.
(Standard states: 1 atm. pressure, 298 K & 1M solution) So, 𝐸𝑟𝑒𝑑𝑢𝑐𝑡𝑖𝑜𝑛
𝜃
= − 𝐸 𝑜𝑥𝑖𝑑𝑎𝑡𝑖𝑜𝑛
𝜃
Example: 𝐸 𝐶𝑢2+/𝐶𝑢
𝜃
= +0.34V so, 𝐸 𝐶𝑢/𝐶𝑢2+
𝜃
= - 0.34V
Reference Electrode
The electrode of known potential is called reference electrode. It may be primary reference electrode like
standard hydrogen electrode (SHE) or secondary reference electrode like calomel electrode.

18. Standard hydrogen electrode (SHE):
• Standard hydrogen electrode (SHE): -
also known as normal hydrogen
electrode (NHE), consists of platinum
wire, carrying platinum foil coated
with finely divided platinum black.
• The wire is sealed into a glass tube.
placed in beaker containing 1 M HCl.
• The hydrogen gas at 1 atm pressure is
bubbled through the solution at 298K.
Half-cell is Pt, H2 (1 atm) /H+ (1 M)

19. Standard hydrogen electrode
• In SHE. at the surface of platinum, either of the following reaction can take place :
• SHE acts as cathode: 2H+(aq) + 2e- → H2(g) (Reduction) → 𝐸 𝐻+/𝐻2
𝜃
OR SHE acts as anode H2(g) → 2H+(aq) + 2e- (Oxidation) → 𝐸 𝐻2/𝐻+
𝜃
• The electrode potential of SHE has been fixed as zero at all temperatures.
• i.e. 𝑬 𝑯+/𝑯 𝟐
𝜽
= 𝑬 𝑯 𝟐/𝑯+
𝜽
= 0.00V
• Its main drawbacks of SHE are
1. It is difficult to maintain 1 atm pressure of H2 gas.
2. It is difficult to maintain H+ ion concentration 1 M.
3. The platinum electrode is easily poisoned by traces of impurities.
• Hence, calomel electrodes are conveniently used as reference electrodes,
• It consists of mercury in contact with Hg2Cl2 (calomel) paste in a saturated solution of KCl.

20. Electromotive Force (e.m.f) of a Cell:
It is the difference between the electrode potentials of two half-cells and cause flow of current from
electrode at higher potential to electrode at lower potential. It is also the measure of free energy change.
Standard emf of a cell is denoted as 𝐸𝑐𝑒𝑙𝑙
𝑜
It is calculated as 𝑬 𝒄𝒆𝒍𝒍
𝜽
= 𝑬 𝒄𝒂𝒕𝒉𝒐𝒅𝒆
𝜽
− 𝑬 𝒂𝒏𝒐𝒅𝒆
𝜽
= 𝑬 𝒓𝒊𝒈𝒉𝒕
𝜽
− 𝑬𝒍𝒆𝒇𝒕
𝜽
where 𝑬 𝒄𝒆𝒍𝒍
𝜽
= Standard e.m.f or Standard cell potential.
(Remember: For numerical purpose always take Eright
θ
& Eleft
θ
values as standard reduction potentials.)
S. No. E.m.f. Cell potential
1 Potential difference between two electrodes
when no current flows in the circuit.
The potential difference of the two half cells when
electric current flows through the cells is called cell
potential
2 Emf is the maximum voltage which can be
obtained from a cell. It gives maximum work
obtained from a cell.
It is always less than the maximum voltage
obtained from the cell.
The work obtained is less than the maximum work
from a cell.
3 It is measured by a potentiometer It is measured by a voltmeter

21. For a Daniell Cell: Zn(s) / Zn2+(aq) // Cu2+(aq) / Cu(s)
Cell reaction: Zn(s) + Cu2+(aq)  Zn2+(aq) + Cu(s)
[Zn acts as anode & Cu acts as cathode]
Standard cell potential : 𝐸𝑐𝑒𝑙𝑙
𝜃
= 𝐸𝑐𝑎𝑡ℎ𝑜𝑑𝑒
𝜃
− 𝐸 𝑎𝑛𝑜𝑑𝑒
𝜃
= 𝐸 𝐶𝑢2+/𝐶𝑢
𝜃
− 𝐸 𝑍𝑛2+/𝑍𝑛
𝜃
𝐸𝑐𝑒𝑙𝑙
𝜃
= + 0.34 – (- 0.76) = 1.10V

22. Electrochemical Series
It is the arrangement of different electrodes in the increasing order of their standard reduction
potentials.
Applications of Electrochemical Series (ECS)
(1)The lower the value of E°redn, the greater the tendency to form cation & the metal undergoes oxidation & is a
reducing agent. M → Mn+ + ne- (Metal, M undergoing oxidation is a reducing agent,
(2) Metals placed below hydrogen in ECS replace hydrogen from dilute acids but metals placed above hydrogen cannot
replace hydrogen from dil acids. Example: Zn + dil. H2SO4  ZnSO4 + H2(g) or
Fe + 2H+  Fe2+ + H2(g).These reactions are feasible as Zn or Fe lie below H2 in ECS.
But, Cu + (dil.) H2SO4  No reaction or Ag + H+  No reaction.
These reaction are not feasible as Cu or Ag lie above H2 in ECS.
(3) Oxides of metals placed below hydrogen are not reduced by H2 but oxides of iron and metals placed above iron are
reduced by H2·
•SnO, PbO, CuO etc are reduced by H2
•CaO, K2O, Na2O etc are not reduced by H2·
(4) Reducing character increases down the series & oxidising character decreasing down the series.
F2 is a best oxidising agent having maximum reduction potential value = + 2.87V
Li is the best reducing agent having least std. reduction value = - 3.05V

23. (5) Reactivity increases down the series (metallic character)
(6) Determination of emf; emf is the difference of reduction potentials of two half-cells.
Eemf = ERHS – ELHS
(a) If Ecell is positive. Then the cell reaction take place spontaneously, otherwise not.(the cell works)
(b) If Ecell < 0, the cell reaction is non spontaneous & the cell does not work.
(c) If Ecell = the cell reaction is in equilibrium.
(7) Greater the reduction potential of a substance, greater is the oxidising power.
(e.g.. F2 > Cl2 > Br2 > I2)
(8) Oxides of metals having E°red ≥ 0.79 V will be decomposed by heating to form O2 and metal.
HgO(s) → Hg(l) + (½)O2(g) Here, (E°Hg
2+
/Hg = 0.79V)
(9) Electrochemical cells are extensively used for determining the pH of solutions, solubility product
(Ksp), equilibrium constant (Kc) and other thermodynamic properties & for potentiometric titrations.
Applications of Electrochemical Series (ECS)
(contd.)
Think: (1) Can we store nickel sulphate solution in a zinc vessel? Take standard reduction values from
the table)
(2) Can we stir ferrous sulphate solution with silver spoon?

25. Nernst Equation
It relates mathematically the electrode potential of an electrode or cell potential of an electrochemical cell
with their standard states at different concentrations & temperature.
The electrode potential of an electrode at any concentration is measured with respect to the SHE.
We take, [solid] = [pure liquid] =1, i.e. [M] = 1.
Molar concentration of ion = [Mn+]
Here, R = 8.314 J/K/mol, T = 298K, n= no. of electrons
taking part in oxidation or reduction process.
1 F = 96500 C
𝟐.𝟑𝟎𝟑𝑹𝑻
𝒏𝑭
=
𝟐.𝟑𝟎𝟑 𝒙 𝟖.𝟑𝟏𝟒 𝒙 𝟐𝟗𝟖
𝒏 𝒙 𝟗𝟔𝟓𝟎𝟎
=
𝟎.𝟎𝟓𝟗
𝒏

26. For a electrochemical cell,
aA + bB  cC + dD where no. of electrons exchanged = n.
Nernst equation can be written as
Ecell = 𝑬 𝒄𝒆𝒍𝒍
𝒐
−
𝟐.𝟑𝟎𝟑 𝑹 𝑻
𝒏𝑭
𝒍𝒐𝒈 𝑸 𝒄 or,
Ecell = 𝑬 𝒄𝒆𝒍𝒍
𝒐
−
𝟐.𝟑𝟎𝟑 𝑹 𝑻
𝒏𝑭
𝒍𝒐𝒈
[𝑪] 𝒄[𝑫] 𝒅
[𝑨] 𝒂[𝑩] 𝒃 where
𝟐.𝟑𝟎𝟑𝑹𝑻
𝒏𝑭
=
𝟐.𝟑𝟎𝟑 𝒙 𝟖.𝟑𝟏𝟒 𝒙 𝟐𝟗𝟖
𝒏 𝒙 𝟗𝟔𝟓𝟎𝟎
=
𝟎.𝟎𝟓𝟗
𝒏
or Ecell = 𝑬 𝒄𝒆𝒍𝒍
𝒐
−
𝟎.𝟎𝟓𝟗
𝒏
𝒍𝒐𝒈
[𝑪] 𝒄[𝑫] 𝒅
[𝑨] 𝒂[𝑩] 𝒃
In general, Ecell = 𝑬 𝒄𝒆𝒍𝒍
𝒐
−
𝟎.𝟎𝟓𝟗
𝒏
𝒍𝒐𝒈 𝑸 𝒄
(where Qc is the reaction quotient and Qc =
[𝑪] 𝒄[𝑫] 𝒅
[𝑨] 𝒂[𝑩] 𝒃 )
NERNST
EQUATION:

27. Example: Write the Nernst equation at 298K for a cell:
Al(s)/Al3+(C1) // Zn2+ (C2) /Zn(s).
Solution:The cell reaction: Anode (Oxidation: Al  Al3+ + 3e [x 2
Cathode (Reduction: Zn2+ + 2e  Zn [ x 3
Cell reaction: 2Al + 3Zn2+ (C2)  2Al3+(C1) + 3Zn
[Here n = 2 x 3 = 6]
The Nernst Equation: Ecell = 𝑬 𝒄𝒆𝒍𝒍
𝒐
−
𝟎.𝟎𝟓𝟗
𝒏
𝒍𝒐𝒈
𝑨𝒍 𝟑+ 𝟐
𝒁𝒏 𝟑
𝑨𝒍 𝟐 𝒁𝒏 𝟐+ 𝟑
= 𝑬 𝒄𝒆𝒍𝒍
𝒐
−
𝟎.𝟎𝟓𝟗
𝟔
log
𝑪 𝟏
𝟐
(𝟏) 𝟑
(𝟏) 𝟐 𝑪 𝟐
𝟑 = 𝑬 𝒄𝒆𝒍𝒍
𝒐
−
𝟎.𝟎𝟓𝟗
𝟔
log
𝑪 𝟏
𝟐
𝑪 𝟐
𝟑 (Ans)
NERNST EQUATION :

28. Nernst equation and Equilibrium constant, Kc
At equilibrium : No net electrons flow in an electrochemical cell,
so Qc = Kc. & Ecell = 0
We know, Ecell = 𝑬 𝒄𝒆𝒍𝒍
𝒐
−
𝟐.𝟑𝟎𝟑 𝑹 𝑻
𝒏𝑭
𝐥𝐨𝐠 𝑸 𝒄 ,
At Equilibrium, 0 = 𝑬 𝒄𝒆𝒍𝒍
𝒐
−
𝟐.𝟑𝟎𝟑 𝑹 𝑻
𝒏𝑭
𝐥𝐨𝐠 𝑲 𝒄
or, 𝑬 𝒄𝒆𝒍𝒍
𝒐
=
𝟐.𝟑𝟎𝟑 𝑹 𝑻
𝒏𝑭
𝐥𝐨𝐠 𝑲 𝒄
Hence, 𝑬 𝒄𝒆𝒍𝒍
𝒐
=
𝟎.𝟎𝟓𝟗
𝒏
𝐥𝐨𝐠 𝑲 𝒄

29. ΔG° is the standard Gibbs free energy change
We know, 𝑬 𝒄𝒆𝒍𝒍
𝒐
=
𝟐.𝟑𝟎𝟑 𝑹 𝑻
𝒏𝑭
𝐥𝐨𝐠 𝑲 𝒄
or, 𝒏𝑭𝑬 𝒄𝒆𝒍𝒍
𝒐
= 𝟐. 𝟑𝟎𝟑𝑹𝑻 𝐥𝐨𝐠 𝑲 𝒄
or, ΔG° = – 2.303 RT log Kc where [ΔG° = - nF𝑬 𝒄𝒆𝒍𝒍
𝒐
]
Relationship between free energy change (ΔG°) and
equilibrium constant (Kc)
Type of reaction ΔG° 𝑬 𝒄𝒆𝒍𝒍
𝒐
Type of cell
Spontaneous -ve +ve Galvanic cell
Non
spontaneous
+ve -ve Electrolytic cell
Equilibrium 0 0 Dead cell

30. Concentration Cell is an electrochemical cell in which two electrodes their ions are same, but
they differ in ionic concentrations of electrolytes. Its 𝑬 𝒄𝒆𝒍𝒍
𝒐
= 𝟎
For a concentration cell: M(s) / Mn+(C1) // Mn+(C2)/M(s)
Here the cell reaction is Mn+ (C2)  Mn+(C1) where no. of electrons = n and C2 > C1
Nernst equation can be written as Ecell = -
𝟐.𝟑𝟎𝟑 𝑹 𝑻
𝒏𝑭
𝒍𝒐𝒈
𝑪 𝟏
𝑪 𝟐
or, Ecell =
𝟐.𝟑𝟎𝟑 𝑹 𝑻
𝒏𝑭
𝒍𝒐𝒈
𝑪 𝟐
𝑪 𝟏
or, Ecell =
𝟎.𝟎𝟓𝟗
𝒏
𝒍𝒐𝒈
𝑪 𝟐
𝑪 𝟏
Case 1: Electrode concentration cell: two hydrogen electrodes or different pressures are dipped in
the same solution of electrolyte, e.g. Pt, H2 (p1) /H+/H2(p2), Pt at 25°C, where p1> p2 and 𝑬 𝒄𝒆𝒍𝒍
𝒐
= 𝟎
Nernst Equation can be written as Ecell =
𝟐.𝟑𝟎𝟑 𝑹 𝑻
𝒏𝑭
𝒍𝒐𝒈
𝒑 𝟐
𝒑 𝟏
or, Ecell =
𝟎.𝟎𝟓𝟗
𝒏
𝒍𝒐𝒈
𝒑 𝟐
𝒑 𝟏
Case 2: Electrolyte concentration cells: Electrodes (say Zn) are the same but electrolyte solutions
have different concentrations, e.g. Zn(s) / Zn2+(C1) // Zn2+(C2)/Zn(s), Here, C2 > C1 & 𝑬 𝒄𝒆𝒍𝒍
𝒐
= 𝟎
Hence, Nernst Eqn. is written as Ecell =
𝟐.𝟑𝟎𝟑 𝑹 𝑻
𝟐𝑭
𝒍𝒐𝒈
𝑪 𝟐
𝑪 𝟏
or, Ecell =
𝟎.𝟎𝟓𝟗
𝟐
𝒍𝒐𝒈
𝑪 𝟐
𝑪 𝟏

31. Numerical problem: Calculate the potential of hydrogen electrode in contact with a solution whose pH is 10.
Solution: We have, H+ (pH = 10) + 1e  ½ H2 , SHE = 0.00V and pH = - log [H+]
Nernst equation: 𝐸 𝐻+/𝐻2
= −
0.059
1
𝑙𝑜𝑔
1
[𝐻+]
= 0.059 (- log [H+]) = 0.059 pH = 0.059 x 10 = 0.59 V (Ans)
Numerical problems for practice:
(1) Write the Nernst equation & calculate the emf of the following cell at 298K: Cu(s)/Cu2+(0.13M)//Ag+(10-4 M)/Ag(s).
(Given Standard reduction potential of Cu & Ag are 0.34V & 0.80V ) [Ans. 0.25V]
(2) A cell contains two hydrogen electrodes. The negative electrode is in contact with a solution of 10-6M hydrogen ions. The emf of the cell is
0.118 V at 298K. Calculate the concentration of hydrogen ions at the positive electrode. Find its pH. [Ans; 10-4M, 4]
(3) Calculate the e.m.f. of the following cell at 298K. 2Cr(s) + 3Fe2+(0.1M)  2Cr3+(0.01M) + Fe(s).
(𝐸 𝐶𝑢2+/𝐶𝑢
𝑜
= +0.34V & 𝐸 𝐹𝑒2+/𝐹𝑒
𝑜
= -0.44 V) [Ans 0.31V]
(4) Consider the following cell reaction: 2Fe(s) + O2 + 4H+(aq)  2 Fe2+(aq) + 2H2O(l) ; 𝐄 𝐜𝐞𝐥𝐥
𝐨
= 1.67V, [Fe2+] =10-3M,
𝑝 𝑂2
= 0.1 atm, The e.m.f. of the cell is 1.57 V. Calculate the pH of the solution. [Ans; pH = 2.95]
(5) Calculate the equilibrium constant (Kc) for the reaction: 4Br- + O2 + 4H+ ⇌ 2Br2 + H2O(l) , Given that 𝐸 𝐶𝑒𝑙𝑙
𝜃
= 0.16 V. Comment on the
reaction. [Ans. 7.03 x 1010, large, the reaction proceeds towards the production of product.]
(6) For a cell reaction at 298 K: A(s) + 2B-(aq) ⇌ A2+ + 2B(s) the equilibrium constant is 104. Calculate 𝑬 𝒄𝒆𝒍𝒍
𝒐
. [Ans. 0.118V]
(7) The zinc/silver oxide cell is used is used in hearing aids and electric watches.
Zn  Zn2+ + 2e (Eo = 0.76V) & Ag2O + H2O + 2e  2Ag + 2OH- (Eo = 0.344 V)
(a) Which is oxidised & which is reduced? (b) Find 𝐄 𝐜𝐞𝐥𝐥
𝐨
𝐚𝐧𝐝 ∆ 𝐫 𝐆 𝛉
. [Ans. 1.104 V, - 2.13 x 102 kJ]
(8) Calculate the e.m.f. of a cell & its standard Gibbs energy at 298K : Al(s)/Al3+(0.01M)//Fe2+(0.022M)/Fe(s). Standard electrode potential of Al
& Fe are -1.66V & -0.44V respectively. [Ans. 1.209 V, - 706.38kJ]