1. Electrochemical Cell

2. A device is used to convert the chemical energy produced in a redox reaction in to Electric energy is called a electrochemical cell or simply a chemical cell.
These are also called as Galvanic cell or simply a Voltaic cell.

3. Reduction Oxidation -loss of electron
Redox Reactions.
Reduction
-gain of electron
-Loss of oxygen
Loss of oxygen
Gain of electrons
Oxidation
-loss of electron
-Gain of oxygen

4. An electrochemical cell is composed to two compartments or half-cells, each composed of an electrode dipped in a solution of electrolyte. These half-cells are designed to contain the oxidation half-reaction and reduction half-reaction separately as shown below.

6. The half-cell, called the anode, is the site at which the oxidation of zinc occurs as shown below.
Zn (s) > Zn+2 (aq) + 2e-
During the oxidation of zinc, the zinc electrode will slowly dissolve to produce zinc ions (Zn+2), which enter into the solution containing Zn+2 (aq) and SO4-2 (aq) ions.

7. The half-cell, called the cathode, is the site at which reduction of copper occurs as shown below.
Cu+2 (aq) + 2e > Cu (s)
When the reduction of copper ions (Cu+2) occurs, copper atoms accumulate on the surface of the solid copper electrode.

8. Salt Bridge
The reaction in each half-cell does not occur unless the two half cells are connected to each other.
It is an inverted U-tube contaning aan electrolyte e.g KCL,KNO3 etc it act as bridge by connecting two half cells
Helps in
To completing the electric circuit
To prevent mixing of soln of two half cell.
To help maintain electric neutrality

10. Electrolytic Cell
It is a device to convert Electric Energy into Chemical Energy
An electrolytic cell is an electrochemical cell in which the energy from an applied voltage is used to drive an otherwise nonspontaneous reaction. Such a cell could be produced by applying a reverse voltage to a voltaic cell like the Daniell cell.

11. Zn(s) +Cu2+(Aq) ---------- Zn2+ (Aq) +Cu(s)
Reaction in Daniell Cell
Zn(s) +Cu2+(Aq)  Zn2+ (Aq) +Cu(s)

12. Electrolytic Cell use full in Electroplating

13. Difference b/w Electrochemical cell and Eletrolytic Cell
Electric enrgy converted into Chemical energy
It is based on redox rxn which is non-spontaneous
Electrochemical cell
Chemical enrgy converted into Electric energy
It is based on redox rxn which is spontaneous i.e rxn occurs its own

14. Electrode Potential
It is the tendency of an electrode in half cell to lose or gain electrons when it is in contact with solution of its own ion
Zn (s) > Zn+2 (aq) + 2e- Oxidation half rxn
Cu+2 (aq) + 2e > Cu (s) Reduction half rxn

16. Electomotive force (emf)
The two half-cells are also connected externally. In this arrangement, electrons provided by the oxidation reaction are forced to travel via an external circuit to the site of the reduction reaction. The fact that the reaction occurs spontaneously once these half cells are connected indicates that there is a difference in potential energy. This difference in potential energy is called an electomotive force (emf) and is measured in terms of volts. The zinc/copper cell has an emf of about 1.1 volts under standard conditions.

17. Nernst Equation
Electrochemistry deals with cell potential as well as energy of chemical reactions. The energy of a chemical system drives the charges to move, and the driving force give rise to the cell potential of a system called galvanic cell. The energy aspect is also related to the chemical equilibrium. All these relationships are tied together in the concept of Nearnst equation.

18. Walther H. Nernst ( ) received the Nobel prize in 1920 "in recognition of his work in thermochemistry". His contribution to chemical thermodynamics led to the well known equation correlating chemical energy and the electric potential of a galvanic cell or battery.

19. The general Nernst equation correlates the Gibb's Free Energy DG and the EMF of a chemical system known as the galvanic cell.
Ecell = E0cell - (RT/nF)lnQ

20. Nernst Equation Remember that G = G + RT ln Q This means
−nFE = −nFE + RT ln Q
Dividing both sides by −nF, we get the Nernst equation:
or, using base-10 logarithms,
E = E − RT nF
ln Q
E = E −
2.303 RT nF
ln Q 20

21. Nernst Equation At room temperature (298 K), and = 0.0592 V
R = J/mol K
F = 96,485 J/V-mol
2.303 RT F
= V
The final form of the Nernst Equation becomes
E = E −
0.0592 n
ln Q

22. Nernst Equation For a general reduction reaction,
The Nernst equation can be written as
(At 298K)
Where n = Number of electrons involved
[Mn+] = molar concentrations at 298K

23. Illustrative Example
Calculate the electrode potential at a copper electrode dipped in a 0.1M solution of copper sulphate at 250C . The standard potential of Cu2+/Cu system is 0.34 volt at 298 K.
Solution:
Cu2+ + 2e-  Cu

24. Calculate the EMF of the cell
Zn(s) | Zn2+ (0.024 M) || Zn2+ (2.4 M) | Zn(s)
Solution
Zn2+ (2.4 M) e = Zn
ReductionZn = Zn2+ (0.024 M) + 2 e
Oxidation----Zn2+ (2.4 M) = Zn2+ (0.024 M), DE° = Net reaction Using the Nernst equation:
(0.024) DE = log (2.4) = (-0.296)(-2.0) = V

25. Reversible Electrochemical Cells
In order for us to make measurements on an electrochemical cell, it must be operating reversibly.
Place an opposing source of potential in the external circuit
Cell operates reversibly and at a constant composition.
we,max = G 25

26. The Measurement of Cell Potentials
Measure the potential of an electrochemical cell when the cell is at equilibrium, i.e., the state between the galvanic and the electrolytic cell.
Counter potential (load) e-
Porous Disk e- e-
Reducing Agent
Oxidizing Agent
Anode
Cathode 26

27. The Work in Transporting Charge
The maximum work
For the passage d electrons from the anode (LHS) to the cathode (RHS)
F = Faraday’s constant = e NA = C/mole 27

28. The Cell Potential
The work to transport charge 28

29. Standard Cell Potentials
From the reaction Gibbs energy
We define 29

30. The Nernst Equation
E represents the standard cell potential, the potential of the cell when all cell components are under standard conditions.
f (all gases) = 1
a (solutes) = 1
T = K
P = 1.00 bar pressure 30

31. Cells at Equilibrium
When the electrochemical cell has reached equilibrium
Kcell = the equilibrium constant for the cell reaction.
Knowing the E° value for the cell, we can estimate the equilibrium constant for the cell reaction. 31

32. Equilibrium Constant Calculations from Cell Potentials
Examine the following cell.
Pt Sn2+ (aq), Sn4+ (aq) Fe3+ (aq) Fe2+ (aq) Pt
Half-cell reactions.
Sn4+ (aq) + 2 e-  Sn2+ (aq) E(Sn4+/Sn2+) = 0.15 V
Fe3+ (aq) + e-  Fe2+ (aq) E (Fe3+/Fe2+) = V
Cell Reaction
Sn2+ (aq) + 2 Fe3+ (aq)  Sn4+ (aq) + 2 Fe2+ (aq)
Ecell = ( V) = 0.62 V 32

33. Standard Reduction Potentials
Standard reduction potentials are intensive properties.
We cannot measure the potential of an individual half-cell!
We assign a particular cell as being our reference cell
Assign values to other electrodes on that basis. 33

34. The Standard Hydrogen Electrode
Eo (H+/H2) half-cell = V e-
f{H2(g)} = 1.00
H2 (g)
a (H+) = 1.00
Pt gauze 34

35. A Galvanic Cell With Zinc and the Standard Hydrogen Electrode.
Porous Disk or
Salt Bridge
Zn(s)
H2 (g)
a(Zn2+) = 1.00
a (H+) = 1.00
Pt gauze
Source of H+ (e.g.,
HCl (aq), H2SO4 (aq))
Zn2+, SO42-
Anode
Cathode 35

36. The Cell Equation for the Zinc-Standard Hydrogen Electrode.
The cell reaction
2 H+ (aq) + Zn (s)  H2 (g) + Zn2+ (aq)
Pt Zn (s) Zn2+ (aq),a=1 H+ (aq), a=1 H2 (g), f=1 Pt
When we measure the potential of this cell Ecell = ERHS - ELHS
but ERHS = E(H+/H2) = V
 Ecell = E(Zn2+/Zn) = V 36

37. The Spontaneous Direction of a Cell Reaction
Examine the magnitude the of the standard cell potential!
If the standard cell potential is positive, the rG is negative! 37

38. The Composition Dependence of the Cell Potential
Nonstandard cell potential (Ecell) will be a function of the activities of the species in the cell reaction.
To calculate Ecell, we must know the cell reaction and the value of Qcell. 38

39. Example For the following system
Pt H2 (g) H+ (aq) Cu2+ (aq) Cu (s) Pt
Calculate the value of the cell potential when the f (H2) = 0.50, a(Cu2+) = 0.20, and a(H+) = 0.40. 39

40. Concentration Cells Electrolyte concentration cell
the electrodes are identical; they simply differ in the concentration of electrolyte in the half-cells. 40

41. Concentration Cells (II)
Electrode concentration cells
the electrodes themselves have different compositions. This may be due to.
Different fugacities of gases involved in electrode reactions (e.g., The H+ (aq)/H2 (g) electrode).
Different compositions of metal amalgams in electrode materials.

42. Applications of Electrochemistry
Measurement of activities and activity coefficients.
Electrochemical series.
Equilibrium constants and thermodynamic functions of cell reactions 42

43. Obtaining Standard Cell Potentials
Look at the following cell
Pt H2 (g) HCl (aq) AgCl (s) Ag (s) Pt
Ecell = E(AgCl/Ag) - E (H+/H2)
= E(AgCl/Ag) 43

44. Ecell Values and Activity Coefficients
In dilute solution, using the DHLL
Plot LHS vs. m1/2
Once Ecell is known, we can obtain experimental estimates of the mean activity coefficients. 44

45. The Calculation of Standard Cell Potentials45

46. Electrochemical Series
Look at the following series of reactions
Cu2+ (aq) + 2 e-  Cu (s) E(Cu2+/Cu) = V
Zn2+ (aq) + 2 e-  Zn (s) E(Zn2+/Zn) = V
Zn has a thermodynamic tendency to reduce Cu2+ (aq)
Pb2+ (aq) + 2 e-  Pb (s) E(Pb2+/Pb) = V
Fe2+ (aq) + 2 e-  Fe (s) E(-Fe2+/Fe) = V
Fe has a thermodynamic tendency to reduce Pb2+ (aq) 46

47. Thermodynamic Information
Note
And 47

48. Activity
When an elecrolyte is dissolved in water the effective concentration of ions in solution is less then actual conc.
The effective conc.of an ion or electrolyte in a solution is called as its Activity represented by “a”
It is related to the actual conc. Expressed as molality “m” by equation
a=  * m

49. Activity Coefficients (g)
 = a/m
Account for deviation from ideal behavior of a solution.
Infinitely dilute solution, no deviations,  = 1
Relatively dilute solutions, deviations from Coulombic (electric) forces of attraction and repulsion  < 1
Concentrated solutions, deviations caused by ionic interactions,  < 1 or  > 1

50. Activity Coefficients
Geometric mean binary activity coefficient
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