Electrochemistry : Electrochemistry By
V.S.Saravanamani
Asst. Prof in Chemistry
Annapoorana Engineering College, Salem
Electrochemistry : Electrochemistry Galvanic (or) Voltaic (or)
Electrochemical cell
It is a device that produces electrical energy at the expense of chemical energy produced in a reaction.
A cell consists of two half cells or electrodes.
A half-cell or electrode contains a metal rod dipped in an electrolytic solution.
Slide 3 : Electrolytic cell: It is a cell in which electrical energy brings about a chemical reaction.
Electrochemical cell: It is a device that produces electrical energy at the expense of chemical energy produced in a reaction.
Slide 4 : Differences between electrolytic cell and electrochemical cell.
Slide 5 : Daniel cell is represented as
Zn ZnSO4 CuSO4 Cu
(IM) (IM)
The Standard emf of the cell is 1.1 volt.
Zn Zn2+ + 2e (at anode)
Cu+2 + 2e Cu (at cathode)
Zn + Cu2+ Zn2+ + Cu (net cell reaction)
Slide 6 : Representation of a galvanic cell
A galvanic cell has two electrodes namely anode and cathode. The following points are kept in view when a cell is represented:
1) Anode is written on the left hand side and cathode on the right side.
Slide 7 : 2) The electrode and the electrolyte are separated by a vertical line to denote the phase boundaries. The physical state is indicated in bracket. For anode, the electrode is written first and then the electrolyte. (e.g)
Pt H2(g) H+ (1M)
(1atm)
Zn ZnSO4 (1M)
For cathode, the electrolyte is written first and then the electrode (e.g.)
CuSO4 (1M) Cu
Slide 8 : 3) Concentrations of electrolytes and pressures of gases are mentioned.
4) Direct contact between the electrolyte solutions is indicated by a single vertical line ().
If they are connected through a salt bridge two vertical parallel lines () are used.
For example, Daniel cell with a salt bridge is represented as follows:
Zn ZnSO4 CuSO4 Cu
(1M) (1M)
5) When emf is mentioned, the temperature is also indicated.
Slide 9 : Reversible cells
Cells obeying three conditions of thermodynamic reversibility are reversible cells.
e.g. Daniel cell
Zn ZnSO4 CuSO4 Cu
(1M) (1M)
When an emf of 1.1 volt is applied from external source, no current flows in the cell. Also no chemical reaction occurs at the electrodes.
When the external emf applied is slightly higher than 1.1 volt, small current flows into the cell. Chemical reactions occur at the electrodes. Zn2+ ions are reduced. Cu atoms are oxidized.
When the external emf applied is slightly less than 1.1 volt, current flows in the opposite direction. Electrode reactions are reversed.
Slide 10 : Irreversible Cells :
Cells not obeying the conditions of thermodynamic reversibility are irreversible cells.
(e.g.) Zinc-silver cell.
Zn | H2SO4(aq) | Ag
The cell reaction is
Zn + H2SO4 ZnSO4 + H2
When external emf slightly greater than the cell emf is applied, the cell reaction is not reversed. This is because H2 has escaped.
Slide 11 : EMF OF A CELL
It is the potential difference that causes flow of electrons from the electrode of higher potential to the electrode of lower potential.
EMF of a cell is related to electrode potentials as follows:
EMF of Cell = [Standard reduction potential of RHS electrode] – [Standard reduction potential of LHS electrode]
W
E cell = E right – E left
Slide 12 : Electrode Potential :
When a metal M is dipped in its salt solution, one of the following reactions occurs depending on the metal :
Positive metal ions pass into the solution :
M Mn+ + ne-(oxidation)
Slide 13 : 2 Positive ions from the solution deposit over the metal.
Mn+ + ne- M(reduction)
When Cu rod is dipped in CuSO4 solution, Cu+2 ions from the solution deposit on metal rod. They attract negative ions from solution. Thus a double layer of ions is formed close to the metal rod. This is called Helmholtz double layer.
Slide 14 : Applications of emf measurement :
Potentiometric titrations are performed.
Solubility of sparingly soluble salt is determined.
Hydrolysis constants of salts are determined.
The valency of ions is determined.
The standard the energy change of a reaction is calculated using the equation
n = Number of electrons involved in the reaction
E = Standard emf of the cell
F = 96500 coulombs.
The equilibrium constant K of a reaction is calculated from Eo.
log K = .
Slide 15 : Factors affecting electrode potential or emf of Cell
nature of the metal
temperature
concentration of metal ions in the solution
Single electrode potential (or) electrode potential :
It is a measure of the tendency of the metal electrode to lose or gain electrons, when it is in contact with its own salt solution. It is developed due to the formation of a double layer around the metal rod.
Standard electrode potential :
It is a measure of the tendency of the metal electrode to lose or gain electrons, when it is in contact with its own salt solution of 1M strength at 25C.
Slide 16 : Measurement of single electrode potential :
Standard hydrogen electrode (SHE) (Primary Reference Electrode)
It has a platinum foil connected to platinum wire and sealed in a glass tube. The platinum foil is dipped in 1M HCl. Hydrogen gas 1 atm pressure is passed through the side arm of glass tube as shown in the figure. The standard electrode potential of SHE is taken as zero. The electrode is represented,
Pt | H2(g) (1 atm) | H+ (1M)
The electrode reaction is
2H+ + 2e- H2
Slide 17 :
Slide 18 : Limitations (or) drawbacks of SHE:
H2 gas reduces many ions like Ag+ and affects compounds of Hg, Ag etc.
It is difficult to get pure H2.
The pressure of H2 is to be kept 1 atm all the time.
It is difficult to set up and transport.
Slide 19 : The electrode potential changes with barometric pressure.
A large volume of test solution is required.
It cannot be used in solutions of redox systems.
The solution may poison platinum surface.
So we use a secondary reference electrode.
Slide 20 : Calomel Electrode (Secondary reference electrode):
It consists of a glass tube containing pure mercury at the bottom. A paste of mercurous chloride covers the mercury. A solution of potassium chloride is present over the paste. The bottom of the tube is sealed with a platinum wire. There is a side tube for electrical contact. The electrode is represented as,
Hg | Hg2Cl2(s) | KCl(aq)
The electrode reaction is,
Hg2Cl2 + 2e- 2Hg + 2Cl-
The electrode potential is, At 25oC,
E = E – 0.0591 log (Cl-)
For saturated KCl, E = +0.242 volt.
Slide 21 :
Slide 22 : Measurement of single electrode potential using a reference electrode (saturated calomel electrode):
The given electrode, say zinc electrode, is coupled with saturated calomel electrode as in the figure. Since the reduction potential of zinc electrode less than that of calomel electrode, zinc acts as anode and calomel as cathode. The cell reaction will be
Zn/ ZnSO4 (1 M) // KCl (satd )/ Hg2Cl2/Hg
Zn + Hg2Cl2 Zn2+ + 2Hg + 2Cl-
The emf of the cell is measured using a potentiometer.
The value of Ecell = 1.002 volt.
Now, Ecell = Eright – Eleft
= Ecal - EZn
1.002 = 0.242 - EZn
EZn = 0.242-1.002
EZn = - 0.76 volt.
Slide 23 : Advantages of Reference Electrode (Calomel Electrode):
Easy to set up.
Easily transportable
Long shelf life
Reproducibility of emf
Low temperature coefficient
Electrode can be used in a variety of solutions.
Eo value is accurately known
Slide 24 : Ion sensitive electrode :
Ion sensitive electrodes have the ability to respond only to a specific ion and develop a potential ignoring other ions in the solution.
Applications of Ion-sensitive electrode :
To determine ions like H+, K+, Li+, etc.
To determine hardness of water (Ca+2 and Mg+2 ions)
To determine concentration of F-, NO3-, CN- etc.
To determine concentration of a gas using gas-sensing electrodes.
To determine pH of a solution using H+ ion sensitive electrode.
The types (classification) of ion-sensitive electrode :
Glass membrane electrodes
Solid state electrode
Pungor or precipitate electrodes
Liquid – liquid electrode
Slide 25 : The types (classification) of ion-sensitive electrode :
Glass membrane electrodes
Solid state electrode
Pungor or precipitate electrodes
Liquid – liquid electrode
Glass Electrode (or) Measurement of pH using glass electrode
Glass electrode contains a thin-walled glass bulb. The glass has low melting point and high electrical conductivity. 0.1M HCl is present in the bulb. A platinum wire is inserted in the acid.
When the glass membrane separates two solutions differing in pH, exchange of H+ ions takes place between the solutions. As a result a potential is developed across the membrane. The potential EG is given by,
EG = EG + 0.0591 pH
Slide 26 :
Slide 27 : Measurement of pH :
The glass electrode is dipped in the given solution. This system is connected to saturated calomel electrode as in the figure. The emf of the resulting cell is measured using a potentiometer.
From the emf, the pH of the solution is calculated as below:
Ecell = Eright – Eleft
Ecell = Ecal – Eglass
Ecell = 0.242 – (EG + 0.0591 pH)
Ecell = 0.242 - EG - 0.0591 pH
pH =
Slide 28 :
Slide 29 : Advantages of Glass Electrode :
It is easily constructed and used
Results are accurate
Electrode is not easily poisoned
Equilibrium is quickly attained
It can be used in strong oxidizing solutions, coloured solutions and in presence of metal ions
Using special glass electrode, pH can be measured from 0 to 12.
It is used in chemical, industrial, biological and agricultural laboratories.
Slide 30 : Disadvantages or Limitations :
Glass has high resistance. So special electronic potentiometer must be used.
It cannot be used in highly alkaline solutions, in pure ethanol or in acetic acid. If the solution pH is more than 12, glass membrane is affected by cations.
Slide 31 : Electrochemical Series :
Electrodes are arranged in the increasing order of their standard reduction potential values. This order is called electrochemical series.
Slide 32 :
Slide 33 :
Slide 34 : Use or Application or Significance of Electrochemical series or emf series:
1. Calculation of Standard emf of a Cell :
We can calculate the standard emf of a cell, by noting the standard reduction potentials from the electrochemical series.
2. Predicting feasibility of a reaction :
The feasibility of a reaction can be predicted from the E value of the corresponding cell reaction.
If is positive, the reaction occurs spontaneously. If is negative, the reaction is not feasible.
Slide 35 : 3. Hydrogen displacement behavior:
We can find out which metals displace hydrogen gas from dilute acids.
Metals with negative electrode potential liberate hydrogen from dilute sulphuric acid.
(e.g.) Zn with E = -0.76 V displaces H2 from dil. H2SO4.
Zn + H2SO4 ZnSO4 + H2
Silver with a positive E value of 0.8, will not displace H2 from dil.H2SO4.
Ag + H2SO4 No reaction
Slide 36 : 4) Determination of Equilibrium constant of a reaction :
Standard electrode potentials are used to determine the equilibrium constants as follows : We know,
- G = nFE
- G = 2.303 RT log K
Hence, 2.303 RT log K = nfE
log K =
Knowing E, n, F, R and T, K can be calculated.
Slide 37 : 5) Displacement of one element by another :
Metals with a lower value of reduction potential will displace metals with a higher reduction potential from their solution.
(e.g.) Zn with E = -0.76V can displace copper (E=+0.34V) or silver (E=+0.8V) from their solution.
Zn + CuSO4 Cu + ZnSO4
Slide 38 : Potentiometric Titrations :
Principle :
We know the potential of an electrode depends on the concentration of the solution in which it is dipped. As the concentration changes, the emf also changes.
In potentiometric titration, we measure the emf of the cell between a reference electrode and indicator electrode. At the end point, there is a drastic change in the potential.
There are three types of potentiometric titrations :
Redox titrations
Precipitation titrations
Acid-base titrations
Slide 39 : Redox Titrations :
Let us consider the titration of FeSO4 versus K2Cr2O7.
FeSO4 solution is taken in a beaker and a platinum electrode (indicator electrode) is dipped in it. It is then connected to a calomel reference electrode. The emf of the resulting cell is measured using a potentiometer.
K2Cr2O7 is added in small quantities from the burette. The Fe2+ concentration decreases because of the reaction.
Fe+2 Fe+3 + e-
Slide 40 :
Slide 41 : So the emf changes as the titration proceeds. At the end point there is a drastic change in the emf.
When the emf is plotted against volume of K2Cr2O7 , we get a curve as in figure. The end point is that point where the slope of the curve is maximum.
A better end point is obtained by plotting E / V against the volume of K2Cr2O7. The resulting curve reaches a maximum at the end point.
Slide 42 : 2. Precipitation Titration :
Let us consider the titration of silver nitrate versus sodium chloride.
Silver nitrate is taken in the beaker and a silver electrode (indicator electrode) is dipped in it. It is then connected to a calomel reference electrode through ammonium nitrate salt bridge. The emf of the resulting cell is measured using a potentiometer.
Slide 43 : NaCl is added in small quantities from the burette. The Ag+ concentration decreases because of the reaction.
Ag+ + Cl- AgCl
So the emf changes as the titration proceeds. At the end point there is a drastic change in the emf.
When the emf is plotted against volume of NaCl, we get a curve as in figure. The end point is that point, where the slope of the curve is maximum.
A better end point is obtained by plotting E / V against the volume of NaCl. The resulting curve reaches a maximum at the end point.
Slide 44 :
Slide 45 : Advantages of Potentiometric Titrations :
The apparatus is cheap and readily available.
Coloured solutions can be titrated.
Very dilute solutions can be accurately titrated.
Fixing of end point is easy.
Several components may be titrated in the same solution.
Indicator is not necessary.
Slide 46 : Conductor :
A substance that permits electricity to flow through it is a conductor. (e.g.) metal, fused salt aqueous solutions of salts, acid base etc.
Non-Conductor :
A substance that does not permit electricity to flow through it is a non-conductor. (e.g.) plastic, wood, many non-metals.
Electrolyte and its conductivity :
Electrolyte like NaCl completely ionizes in solution.
NaCl Na+ + Cl-
The ions formed move in an electric field to oppositely charged electrodes and conduct electricity.
Slide 47 : Strong electrolyte :
A strong electrolyte is completely ionized in solution at all concentrations. E.g. NaCl.
NaCl Na+ + Cl- (100%)
Weak electrolyte :
A weak electrolyte is partially ionized in solution. (e.g.) CH3COOH.
CH3COOH CH3COO- + H+ (Partial)
Slide 48 : Conductometric Titrations :
The conductance of a solution depends on the number of ions, nature and charge of the ion and its mobility. During a titration, there is a change in the number and nature of ions in solution. Hence there is a change in conductance. This can be used to detect end point of a titration. This type of titration is called conductometric titrations. The temperature should be maintained constant throughout. The titrant should be 10 times stronger than the solution to be titrated so that volume change is very small.
Slide 49 : Titration of strong acid versus strong base :
Let us consider the titration of HCl versus NaOH.
A known volume of HCl is taken in a beaker. The conductivity cell is dipped in the acid. The NaOH is added from the burette in small volumes and the conductance is measured each time. When the alkali is added, the fast moving H+ ions is replaced by slow moving Na+ ions. So the conductance decreases until all the acid is neutralized.
H+ + Cl- + (Na+ + OH-) Na+ + Cl- + H2O
After neutralization, further addition of NaOH increases the conductance sharply due to the presence of fast moving OH- ions in the solution.
Slide 50 : A graph is plotted between conductance and volume of NaOH added. The point of intersection of two straight lines gives the end point.
Slide 51 : Titration of weak acid versus strong base:
Let us consider the titration of acetic acid versus NaOH.
A known volume of acetic acid is taken in a beaker. The conductivity cell is dipped in the acid. The NaOH is added from the burette in small volumes and the conductance is measured each time. When the alkali is added, the conductance of the solution increases due to the formation of completely ionized sodium acetate.
CH3COOH + (Na+ + OH-) CH3COO- + Na+ + H2O
After neutralization, further addition of NaOH increases conductance sharply due to the presence of fast moving OH- ions in the solution.
Slide 52 : A graph is plotted between conductance and volume of NaOH. The point of intersection of two straight lines gives the end point.
Slide 53 : Titration of a mixture of
weak and strong acids versus strong base:
Let us consider the titration of a mixture of HCl and CH3COOH versus NaOH.
A known volume of HCl and CH3COOH is taken in a beaker. The conductivity cell is dipped in the acid. The NaOH is added from the burette in small volumes and the conductance is measured each time.
Slide 54 : First the strong acid HCl is neutralized and the conductance decreases until all the acid is neutralized. Then the neutralization of CH3COOH takes place. The conductance slowly increases, until all CH3COOH is neutralized. Further addition of alkali increases the conductance
A graph is plotted between conductance and volume of NaOH. The first end point corresponds to neutralization of HCl. The second end point corresponds to neutralization of CH3COOH.
Slide 55 :
Slide 56 : Advantages (Merits) of Conductometric titrations :
Indicator is not necessary.
Dilute solutions can be titrated.
Accurate end point is obtained.
No special attention is necessary near the end point.
Coloured solutions can be titrated.
Weak acid can be titrated against weak base.
Slide 57 : Nernst equation for electrode potential :
Let us consider the reaction
Mn+ + ne- M ----- (1)
The free energy change of this equilibrium G is related to the equilibrium constant K by the
Vant Hoff isotherm.
We know, G = - nFE
G = -nFE ------ (3)
Substituting equations (1) & (3) in equation (2)
Dividing by –nF and using the fact that activity of solid metal [M]=1, we have
At 25C, R = 8.314J/K/mol , F = 96500 coulombs. So equation (5) becomes,
Slide 58 : Applications of Nernst Equation :
To calculate electrode potential of unknown metal.
To predict corrosion of metals.
To set up electrochemical series.
Nernst equation for a reversible cell :
Let us consider the reaction in a reversible cell :
A + B C + D ----- (1)
The free energy change G of this equilibrium is related to the equilibrium constant K by the Van’t Hoff isotherm.
We know, G = - nFE
G = -nFE ------ (3)
Substituting equations (1) & (3) in equation (2)
Slide 59 : Dividing by –nF
This is Nernst equation for a reversible cell. At 25C, R = 8.314J/K/mol,
F = 96500 coulombs. So equation (5) becomes,
Slide 60 : THE END