CLASS 12 ELECTROCHEMISTRY.pptx

Electrochemistry notes with Explanation

1. SUB CHEMSITRY CLASS 12 ELECTROCHEMSITRY

2. 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. 2. Electroplating. 3. Purification of metals. 4. Batteries and cells used in various instruments.

3. Conductors Substances that allow electric current to pass through them are known as conductors. 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. 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. 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 , K 2 SO 4 2. Weak electrolytes The electrolytes that dissociate partially are called weak electrolytes, e.g., CH 3 COOH, H 2 CO 3 , NH 4 OH,H 2 S, etc.

4. ELECTROCHEMICAL CELL The cell which converts chemical energy to electrical energy and vice versa is called an electrochemical cell. TYPES Galvanic Cells -Converts chemical energy into electrical energy. Galvanic cell is also called voltaic cell. Electrolytic Cells -Converts electrical energy into chemical energy.

5. General Representation of an Electrochemical Cell CELL DIAGRAM OR REPRESENTATION OF A CELL

6. example The Daniel cell is represented as follows : Zn(s) | Zn2+ (C1 ) || Cu2+ (C2 ) | Cu (s) Salt bridge When the oxidation takes place then it is anode when there is reduction takes place then thereis cathode.

7. GALVANIC CELL DANIEL CELL It is a Galvanic Cell in which Zinc and Copper are used for the redox reaction to take place. Zn (s) + Cu2+ ( aq ) Zn2+ ( aq ) + Cu(s) Oxidation Half : Zn (s) Zn2+ ( aq ) + 2e– Reduction Half : Cu2+( aq ) + 2e– Cu(s) NOTES:- Anode is assigned negative polarity cathode is assigned positive polarity. In Daniel Cell, Zn acts as the anode and Cu acts as the cathode Ques …. What is the difference between electrolytic cell and galvanic cell.

8. GALVANIC CELL

9. FUNCTION OF SALT BRIDGE It completes the circuit and allows the flow of current. It maintains the electrical neutrality on both sides. Salt-bridge generally contains solution of strong electrolyte such as KNO 3 , KCL etc. KCl is preferred because the transport numbers of K +  and Cl -are almost same.

10. SOME TERMS CELL POTENTIAL It is the difference between  electrode potentials  (of both  electrodes   i.e anode and cathode) of the given cell. It is denoted by E cell. Electrode Potential 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. OR Potential difference between the electrode and electrolyte. Oxidation potential The tendency to lose electrons in the above case is known as oxidation potential. Oxidation potential of a half-cell is inversely proportional to the concentration of ions in the solution. Reduction potential The tendency to gain electrons in the above case is known as reduction potential. 

11. E cell is written as : E cell = E right – Eleft EXAMPLE:- Cell reaction: Cu(s) + 2Ag+ ( aq ) → Cu 2+( aq ) + 2 Ag(s) Half-cell reactions: Cathode (reduction): 2Ag+ ( aq ) + 2e– → 2Ag(s) Anode (oxidation): Cu(s) → Cu 2+( aq ) + 2e The cell can be represented as: Cu(s)|Cu 2+( aq )||Ag+ ( aq )|Ag(s) Ecell = E right – E left = E Ag+/ Ag – E Cu2+ /Cu NOTES:- It is not possible to determine the absolute value of electrode potential. For this a reference electrode [NHE or SHE] is required.

12. STANDARD HYDROGEN ELCTRODE(SHE) It is related to standard conditions. The standard conditions taken are : 1M concentration of each ion in the solution. (ii) A temperature of 298 K. (iii) 1 bar pressure for each gas

13. EMF(ELECTROMOTIVE FORCE) 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. It is denoted by E0( prnounced as e not).

14. ELECTROCHEMICAL SERIES The half cell potential values are standard values and are represented as the standard reduction potential values as shown in the table at the end which is also called Electrochemical Series.

15. APPLICATIONS OF ELECTROCHEMICAL SERIES The lower the value of E°, the greater the tendency to form cation . M → M n +  + ne Reducing character increases down the series. Reactivity increases down the series. Determination of emf ; emf is the difference of reduction potentials of two half-cells. • E emf  = E RHS  – E LHS If the value of emf is positive. then reaction take place spontaneously, otherwise not.

16. NERNST EQUATION It relates electrode potential with the concentration of ions. .

17. NERNST EQUATION Concentration of pure solids and liquids is taken as unity.

18. APPLICATIONS OF NERNST EQUATION Relationship between free energy change and equilibrium constant ΔG° = – 2.303RT log K c

19. CONDITIONS FOR SPONTANEOUS REACTION( IMPORTANT)

20. CONDUCTANCE: It is the reciprocal of resistance. T he ease with which the electric current flows through a conductor.It is denoted by G. G = (1/R), units ohm -1  mhos or Ω -1 SPECIFIC CONDUCTIVITY -It is the reciprocal of resistivity .It is denoted by kappa(k).

21. NOTE:-Specific conductivity decreases on dilution. This is because concentration of ions per cc decreases upon dilution. Molar Conductivity ( Λm ) The conductivity of all the ions produced when 1 mole of an electrolyte is dissolved in V mL of solution is known as molar conductivity.It is related to specific conductance as Λ m  = (k x 1000/M) where. M = molarity . It units are Ω -1  cm 2  mol -1  or S cm 2  mol -1 . Equivalent conductivity ( Λ e ) The conducting power of all the ions produced when 1 g-equivalent of an electrolyte is dissolved in V mL of solution, is called equivalent conductivity. It is related to specific conductance as Λ e  = (k x 1000/N) where. N = normality. Its units are ohm -1  cm 2  (equiv -1 ) or mho cm 2  (equiv -1 ) or S cm 2  (g-equiv -1 ).

22. MOLAR CONDUCTIVITY & EQUIVALENT CONDUCTIVITY

23. FACTORS AFFECTING ELECTROLYTIC CONDUCTANCE Electrolyte -is a substance that dissociates in solution to produce ions and hence conducts electricity in dissolved or molten state. Examples : HCl , NaOH , KCl (Strong electrolytes). CH 3 COOH, NH 4 OH (Weak electrolytes). Nature of electrolyte or interionic attractions Temperature

24. VARIATION OF CONDUCTIVITY AND MOLAR CONDUCTIVITY WITH DILUTION Conductivity always decreases with decrease in concentration both, for weak and strong electrolytes. This can be explained by the fact that the number of ions per unit volume that carry the current in a solution decreases on dilution. Molar conductivity increases with decrease in concentration. This is because the total volume, V, of solution containing one mole of electrolyte also increases

25. LIMITING MOLAR CONDUCTIVITY (λ0m ) The value of molar conductivity when the concentration approaches zero is known as limiting molar conductivity or molar conductivity at infinite dilution. It is possible to determine the molar conductivity at infinite dilution λ0m. In case of strong electrolyte by extrapolation of curve of λ0mvs c. In caseof weak electrolyte at infinite dilution cannot be determined by extapolation of the curve as the curve becomes almost parallel to y-axis when concentration approaches to zero.

26. KOHLRAUSCH’S LAW It states that the limiting molar conductivity of an electrolyte can be represented as the sum of the individual contributions of the anion and cation of the electrolyte. In general, if an electrolyte on dissociation gives v+ cations and v– anions then its limiting molar conductivity is given by:-

27. APPLICATIONS OF LAW

28. ELECTROLYSIS It is the process of decomposition of an electrolyte when electric current is passed through either its aqueous solution or molten state. PREDICT THE PRODUCT OF ELECTROLYSIS:- when an aqueous solution of an electrolyte is electrolysed , if the cation has higher reduction potential than water (-0.83 V), cation is liberated at the cathode (e.g.. in the electrolysis of copper and silver salts) otherwise H 2 gas is liberated due to reduction of water (e.g., in the electrolysis of K, Na, Ca salts, etc.) Similarly if anion has higher oxidation potential than water (- 1.23 V), anion is liberated (e.g., Br - ), otherwise O 2  gas is liberated due to oxidation of water (e.g., in caseof F-, aqueous solution of Na 2 SO 4  as oxidation potential of SO 2- 4  is – 0.2 V).

29. BATTERIES Cells and batteries are available in a wide variety of types. TYPES A primary cell cannot be recharged and cannot be used. A secondary cell, or storage cell, can be recharged and can be use.

30. PRIMARY CELL DRY CELL:- Anode : Zn container Cathode : Carbon (graphite) rod surrounded by powdered MnO 2 and carbon. Electrolyte : NH 4 Cl and ZnCl 2 Reaction : Anode : Zn → Zn 2+ + 2e– Cathode : MnO 2 + NH 4 + + e – → MnO (OH) + NH 3 The standard potential of this cell is 1.5 V and it falls as the cell gets discharged continuously and once used it cannot be recharged.

31. MERCURY CELL: It offers a rather more stable voltage. The emf of the Mercury Cell is 1.35 V. Usually, the mercury cell is costlier. This is the reason, why they are used only in sophisticated instruments such as camera, hearing aids, and watches etc.

32. The reactions during discharge are, At anode: Zn(Hg) + 2OH –  →Zn (OH) 2  + 2e – At cathode: HgO + H2O + 2e –  →Hg + 2OH – Overall reaction: Zn(Hg) + HgO (s) →Zn(OH) 2 + Hg(l)

33. SECONDARY CELL Secondary cells are those which can be recharged again and again for multiple uses. e.g. lead storage battery and Ni – Cd battery.

34. LEAD STORAGE Anode : Lead ( Pb ) Cathode : Grid of lead packed with lead oxide (PbO 2 ) Electrolyte : 38% solution of H 2 SO 4 Discharging Reactions Anode: Pb (s) + SO 4 2–( aq ) PbSO 4 (s) + 2e– Cathode: PbO 2 (s) + 4H + ( aq ) + SO4 2–( aq ) + 2e– PbSO 4 (s) + 2H 2 O(l) Overall Reaction : Pb (s) + PbO 2 (s) + 2H 2 SO 4 ( aq ) 2PbSO 4 (s) + 2H 2 O(l)

35. CORROSION It involves a redox reaction and formation of an electrochemical cell on the surface of iron or any other metal. Anode : 2Fe (s ) → 2 Fe 2+ + 4e– Eº = + 0.44 V Cathode : O2 (g) + 4H+ + 4e– → 2H2 O(l) Eº = 1.23 V Overall reaction: 2Fe (s) + O 2 (q) + 4H + → 2Fe 2+ + 2H2O

36. Important questions a) Following reactions occur at cathode during the electrolysis of aqueous silver chloride solution : Ag+( aq ) + e– → Ag(s) E° = +0.80 V H+( aq ) + e– →1/2 H2 (g) E° = 0.00 V On the basis of their standard reduction electrode potential (E°) values, which reaction is feasible at the cathode and why ? (b) Define limiting molar conductivity. Why conductivity of an electrolyte solution decreases with the decrease in concentration ? (c) Calculate emf of the following cell at 25 °C : Fe | Fe2+(0.001 M) || H+(0.01 M) | H2(g) (1 bar) | Pt(s) E°(Fe2+ | Fe) = –0.44 V E°(H+ | H2) = 0.00 V (d) From the given cells : Lead storage cell, Mercury cell, Fuel cell and Dry cell . Answer the following : ( i ) Which cell is used in hearing aids ? (ii) Which cell was used in Apollo Space Programme ? (iii) Which cell is used in automobiles and inverters ? (iv) Which cell does not have long life ?

37. (e)Calculate the degree of dissociation (α) of acetic acid if its molar conductivity (∧m) is 39.05 S cm2mol–1. Given λo (H+) = 349.6 S cm2 mol–1 and λo (CH3COO–) = 40.9 S cm2mol–1. (f) Calculate the mass of Ag deposited at cathode when a current of 2 amperes was passed through a solution of AgNO3 for 15 minutes. (Given : Molar mass of Ag = 108 g mol–1 1F = 96500 C mol–1) (g) Define fuel cell. (h) Write the cell reaction and calculate the e.m.f . of the following cell at 298 K : Sn (s) | Sn2+ (0·004 M) || H+(0·020 M) | H2(g) (1 bar) | Pt (s) (Given :E0 Sn2+/ Sn = – 0·14 V) ( i ) Give reasons : On the basis of Eo values, O2 gas should be liberated at anode but it is Cl2 gas which is liberated in the electrolysis of aqueous NaCl . Conductivity of CH3COOH decreases on dilution. (j) For the reaction 2AgCl (s) + H2 (g) (1 atm ) 2Ag (s) + 2H+ (0·1 M) + 2Cl–(0·1 M), ∆Go= – 43600 J at 25 C. Calculate the e.m.f . of the cell. [log 10–n= – n] (k) Define fuel cell and write its two advantages.

38. (l) E°cell for the given redox reaction is 2.71 V Mg+ Cu2+(0.01v)  Mg2+(0.001V) + Cu (s) Calculate Ecell for the reaction. Write the direction of flow of current when an external opposite potential applied is i ) less than 2.71 V and (ii) greater than 2.71 V (m) A steady current of 2 amperes was passed through two electrolytic cells X and Y connected in series containing electrolytes FeSO4 and ZnSO4 until 2.8 g of Fe deposited at the cathode of cell X. How long did the current flow? Calculate the mass of Zn deposited at the cathode of cell Y. (Molar mass : Fe = 56 g mol–1 Zn = 65.3 g mol–1, 1F = 96500 C mol–1)

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