Electrochemistry

1. Chemistry (1)
CHE 011
Preparatory year
Lecture (2)
Electrochemistry
Ass. Prof./ Heba Abd El Gawad
The Higher Institute of Engineering
Chemical Engineering Department
2022-2023

2. Definition of electrochemistry:
 It is defined as the branch of chemistry which deals with interconversion of chemical
energy to electrical energy and vice versa.
 For example:
I. In battery and full cell, chemical energy is converted to electrical energy.
Spontaneous
I. In electroplating/electrolysis and battery charging, electrical energy is converted to
chemical energy. Non-Spontaneous

3. Electrochemical cell:
 It is a system or device for the conversion of electrical energy into chemical energy or
vice-versa.
 In which two electrodes are fitted in the same electrolyte or in two different
electrolytes, which are joined by a salt bridge.
 Types of electrochemical cell:
1. Galvanic cell or voltaic cell
The device used to convert the chemical energy produced on a redox reaction into
electrical energy in the form of an electrical current.
Chemical energy Electrical energy
Spontaneous

4. 2. Electrolytic cell
It is a device in which electrolysis (chemical reaction involving oxidation and
reduction) is carried out by using electricity or in which the conversion of electrical
energy into chemical energy.
Electrical energy Chemical energy
Non-Spontaneous

5. Redox reaction:
It is a chemical reaction in which electrons are transferred from one specie to another.
A species losing electrons is said to be oxidized; one gain electrons is said to be
reduced. The two processes together are called redox (ie, oxidation and reduction
reactions can occur).
Oxidation Is the loss of electrons
(Increase in oxidation number) Redox reaction
Reduction Is the gain of electrons
(Decrease in oxidation number)

6. Oxidation reaction
 Example:
Reduction reaction
Redox reaction
Oxidizing agent
Reducing agent

7. Galvanic cell
 A typical example of a galvanic cell is that of a Daniell cell.
 Daniell cell is a galvanic cell in which Zinc (Zn) and Copper (Cu) are used
for the redox reaction to take place.
 It is formed by combination of two half cells. One is oxidation half cell or
anodic half cell and other one is reduction half cell or cathodic half cell.

8.  The oxidation half cell contains a zinc metal electrode dipping in ZnSO4
solution and the reduction half cell consists of Cu metal electrode in
CuSO4 solution.
 Both the half cells are connected externally by a metallic wire and internally
connected by a salt bridge.
 A salt bridge is a U shaped tube containing a salt solution such as KCl,
KNO3, Na2SO4 or NH4NO3 that prevents the mechanical mixing of the
solution joins the half cells.

10.  The following reactions take place in the cell:
At anode: Zn Zn2+ + 2e- (Oxidation reaction)
At cathode: Cu+2 + 2e- Cu (Reduction reaction)
The overall cell reaction:
Zn + Cu+2 Zn2+ + Cu (Redox reaction)
Presentation of galvanic cell:
Zn Zn2+(C1) Cu+2(C2) Cu

11.  In galvanic cell, two electrodes are represented; at one electrode oxidation takes place
and at other electrode reduction takes place.
 The tendency of an electrode (M) to get oxidized or reduced when it is contact with its
own ionic (Mn+) solution, known as electrode potential. It is represented by E.
 If the metal electrode is suspended in a ionic solution 1M (molar) concentration or
unity activity and temperature is kept at 25 oC under 1 atm pressure, the potential
developed between electrode and ionic solution is known as standard electrode
potential. It represented by Eo.
Electrode potential

13.  Potential difference set up between metal rod and solution containing the metal
ions, and depends upon metal nature and metal ion concentration.
 The electrical potential difference between the cathode and the anode is known as:
1. Cell potential: is known as reduction potential (E)
2. Electromotive force (EMF)

14. Electromotive force (EMF)
 The force which causes the flow of electrons from one electrode to another
electrode and thus results in the flow of current from electrode at higher
potential to electrode at lower potential is called electromotive force.
Standard EMF or standard cell potential = (Standard reduction potential of R.H.S. electrode –
(Standard reduction potential of L.H.S. electrode)
Oxidation potential = -reduction potential

15. Electrochemical series
 Series in which elements are arranged in increasing order of their standard electrode
potential is known as electrochemical series.
 The standard electrode potential for any given element to be oxidized or
reduced in a galvanic cell can be measured by coupling it with a reference
electrode.
 Standard hydrogen electrode is the reference electrode, .

17. Example:

19. Example:

20. Nernst equation
Ecell = Eo - RT/nF Ln [K]
 Where:
E is the electrode potential
Eo is the standard electrode potential
R is the universal gas constant
n is the number of electrons
F is the faraday (96,500 coulombs)
K is equilibrium constant
K = concentration of product / concentration of reactant

21. mMn+ nNn+
[Nn+]n
P
K =
[Mn+]m
R
Ecell= Eo - 0.059/n Log 𝑲
This equation used to measure the cell potential at temperature
not equal 25 oC and concentration of electrolyte not equal 1
molar. (non standard conditions).

22. Nernst equation can be used to calculate:
1. Half-cell potential
2. Cell potential

23. Example:

25. Spontaneity of cell reaction
 If Ecell = + ve and ∆ G < zero (-ve) reaction is spontaneous
 If Ecell = - ve and ∆ G > zero (+ve) reaction is non-spontaneous
 If Ecell = zero and ∆ G = zero no reaction (reach to equilibrium)
Eo = 0.059/n Log 𝑲
Where: ∆ G = Free energy