2-Electrochemistry.pptx

Reference electrodes and its construction

1. CHAPTER 1: ELECTROCHEMISTRY CHEMISTRY- MODULE I- B.Tech I year Dr. B. Divya Assistant Professor Department of Chemistry Institute of Aeronautical Engineering (IARE) Hyderabad Outcomes: Understand the reference electrodes and its uses Electrochemical series and its applications

2. Don’t write or place any image in this area ELECTROCHEMISTRY TOPICS Reference electrodes Construction of Standard Hydrogen Electrode Secondary Reference Electrode Construction of Calomel Electrode Electrochemical Series & its applications Nernst Equation and its applications

3. Don’t write or place any image in this area ELECTROCHEMISTRY Reference Electrodes The electrode potential is found out by coupling the electrode with primary reference electrode and the potential of which is arbitrarily fixed as zero . The important primary reference electrode used is the Standard Hydrogen Electrode (SHE) Standard Hydrogen Electrode Consists of Pt wire in an inverted glass tube H 2 gas is passed through tube at 1 atm Pt foil is attached at end of wire The electrode is immersed in 1M H + ion at 25°C The electrode potential is Zero at all temperatures

4. Don’t write or place any image in this area ELECTROCHEMISTRY Cell Representation of SHE It is represented as Pt, H 2 (1atm)/H + (1M) If the potential of the coupled electrode is greater than zero then the standard hydrogen electrode acts as anode , the electrode reaction can be written as H 2(g) 2H + +2e - If the potential of the coupled electrode is less than zero then standard hydrogen electrode acts as cathode , the electrode reaction can be written as 2H + +2e - H 2(g) Based on the electrode potential obtained with reference to hydrogen, electrochemical series is obtained

5. Don’t write or place any image in this area ELECTROCHEMISTRY Limitations of SHE SHE is difficult to construct . Bubbling H 2 gas at 1 atm around Pt through the solution is difficult. Maintaining activity of H + at unity is difficult. Pt metal has high adsorption and catalytical activity. Hydrogen on Pt surface is reduction catalyst. A number of organic substances are reduced. Inorganic ions like Fe 3+ , MnO 4- etc reduce to their lower valence state. Cations of Ag, Cu, Cd, Pb will get reduce and deposit on Pt metal . Substance like arsenic and sulfides also interfere.

6. Don’t write or place any image in this area ELECTROCHEMISTRY Secondary Reference Electrode – its need The use of SHE is difficult, because it is difficult to maintain 1M H + ion concentration and the pressure of the gas at one atmosphere. Also, the electrode will easily get poisoned in case of traces of impurities in the gas and hence, other reference electrodes are used Example: Saturated calomel electrode (saturated KCl ) Physical appearance of calomel electrode

7. Don’t write or place any image in this area ELECTROCHEMISTRY Calomel Electrode The saturated calomel electrode is a reference electrode based on the reaction between the elemental mercury and mercury (I) chloride. The aqueous phase in contact with the mercury and mercury (I) chloride (Hg 2 Cl 2 “calomel”) is the saturated KCl in water. In cell notation the electrode is written as: Cl - / Hg 2 Cl 2 (s)/Hg(l)/Pt It consists of glass tube , that contains Hg at the bottom covered with solid Hg 2 Cl 2 Tube is filled with KCl . A platinum wire is in touch with Hg and act as electrical contact

8. Don’t write or place any image in this area ELECTROCHEMISTRY Calomel Electrode-reaction The electrode reaction (reduction),   The potential of the electrode is given by, ) =   )   Calomel electrode is convenient to handle and it can be constructed easily. It cannot be used above 50 ⁰C due to instability of Hg 2 Cl 2 and mercury is poisonous. 0.1 M KCl /Hg 2 Cl 2(s) / Hg,Pt 0.33V 1 M KCl /Hg 2 Cl 2(s) /Hg, Pt 0.28 V Sat KCl / Hg 2 Cl 2(s) /Hg, Pt 0.24 V Electrode potentials at different concentrations of Chloride ions

9. Don’t write or place any image in this area ELECTROCHEMISTRY Electrochemical Series The arrangement of different electrode potentials of different electrodes from highest – ve to highest + ve is called electrochemical series - ve SRP + ve SRP - ve sign indicates that reduced form lose electrons and get oxidized easily

10. Don’t write or place any image in this area ELECTROCHEMISTRY Applications of Electrochemical Series Calculation of standard EMF of the cell Relative ease of oxidation or reduction Displacement of one element by the other Predicting spontaneity or feasibility of redox reaction Hydrogen displacement behavior

11. Don’t write or place any image in this area ELECTROCHEMISTRY Applications of Electrochemical Series  EMF of the electrochemical cell can be calculated by taking the values of electrode potentials of the two half – cells. There are usually three methods that can be used for the calculation: By taking into account the oxidation potential of anode and reduction potential of cathode . By considering the reduction potentials of both electrodes . By taking the oxidation potentials of both electrodes . Calculation of standard EMF of the cell Example: The cell potential of a galvanic cell = Potential of the half cell cathode – Potential of the half cell anode =+0.34V– (-0.76 V) = 1.1 V

12. Don’t write or place any image in this area ELECTROCHEMISTRY Metals at the top of the series are good at giving away electrons. They are good reducing agents. The reducing ability of the metal increases as you go up the series.  Relative ease of oxidation or reduction Metal ions at the bottom of the series are good at picking up electrons. They are good oxidising agents. The oxidising ability of the metal ions increases as you go down the series. The more negative the E° value, the more the position of equilibrium lies to the left Eg : Li Li + + 2e -

13. Don’t write or place any image in this area ELECTROCHEMISTRY Applications of Electrochemical Series Displacement reaction: Displacement of metals Any metal placed above in the electrochemical series displace the metal in the below series in its aqueous solution Fe 2+ /Fe -0.44 V Cu 2+ /Cu +0.34 V Fe (oxidation) (reduction)   ( overall )   Fe/Fe 2+ +0.44 V Cu 2+ /Cu +0.34 V

14. Don’t write or place any image in this area ELECTROCHEMISTRY Applications of Electrochemical Series Displacement of non-metal by other non-metals Reactivity of non-metal increases f rom top to bottom. Non-metal placed below in the series displaces the non-metal from its solution placed above Example: Chlorine displaces iodine from KI solution, because chlorine is placed below iodine in the electrochemical series (oxidation)    

15. Don’t write or place any image in this area ELECTROCHEMISTRY Calculation of standard EMF of a cell and determination of cell reaction (spontaneity) Fe / // / E Fe2+/Fe = -0.44 V E Cu2+/Cu = +0.34 V E cell = E right –E left (Fe dipped in Cu solution) =E Cu2+/Cu – E Fe2+/Fe = 0.34 – (-0.44) = 0.78 V   E is + ve , spontaneous // / Fe   E Fe2+/Fe = -0.44 V E Cu/Cu2+ = -0.34 V E cell = E right –E left =E Cu/Cu2+ – E Fe2+/Fe = -0.34 – (-0.44) = - 0.78 V E is - ve , non- spontaneous ( Cu dipped in Fe solution)

16. Don’t write or place any image in this area ELECTROCHEMISTRY Determination of standard free energy and equilibrium for the reaction We know that, Δ G = Δ G + RT ln Q At equilibrium, Δ G =0 - Δ G = RTln k = 2.303RT log k = nFE   From the value of E , the equilibrium constant for the cell reaction can be calculated.

17. Don’t write or place any image in this area ELECTROCHEMISTRY Equilibrium constant K Equilibrium Products Reversible Reaction Reactants time The change in free energy is equal to zero at equilibrium The equilibrium constant of a chemical reaction (usually denoted by the symbol K) provides insight into the relationship between the products and reactants when a chemical reaction reaches equilibrium At a particular temperature, the rate constants are constant. The ratio of the rate constant of forward reaction to the rate constant of backward reaction should be a constant and is called an equilibrium constant ( Kequ ). Equilibrium Constant Formula K equ  = k f /k b  = [C] c  [D] d /[A] a  [B] b  = K c (or) K equ where Kc, indicates the equilibrium constant measured in moles per litre .   aA + bB ⇒ cC + dD

18. Don’t write or place any image in this area ELECTROCHEMISTRY Applications of Equilibrium constant Predicting the Extent of Reaction: The equilibrium constant (K c ) can be used to predict the extent of a reaction, i.e. the degree of the disappearance of the reactants. The magnitude of the equilibrium constant gives an idea of the relative amount of the reactants and the products. aA + bB ⇒ cC + dD L arger value of the equilibrium constant (>10 3 ) shows that forward reaction is favored H 2 (g) + 12O 2 (g) ⇌ H 2 O(g) ⇒ Kc = 2.4×10 47 Intermediate value of equilibrium constant (10 -3  to 10 3 ) show that the concentration of the reactants and products are comparable. H 2  (g) + I 2  (g) ⇌ 2HI (g) ⇒ Kc = 57 at 700 K .  Low value of equilibrium constant (<10 -3 ) shows that backward reaction is favored H 2 O (g) ⇌ H 2  (g) + (1/2) O 2  (g) ⇒ Kc = 4.1 × 10 -48

19. Don’t write or place any image in this area ELECTROCHEMISTRY Applications of Equilibrium constant Equilibrium Constant for Predicting the Direction of a Reaction The equilibrium constant can be used to predict the direction of the reaction. We need a term, reaction quotient (Qc expressed in terms of concentrations or Qp in terms of partial pressures) similar to the equilibrium constant except that the conditions are not at equilibrium. For a  balanced reaction , aA + bB ⇌ cC + dD Reaction quotient (Qc or Qp ) is given as: Q c  = [C] c [D] d /[A] a [B] b

20. Don’t write or place any image in this area ELECTROCHEMISTRY Hydrogen Displacement Behavior Example: Zinc and iron lie above hydrogen in reactivity series and can displace hydrogen from acid. Zn + H 2 SO 4  → ZnSO 4  + H 2  ↑ (dilute) Fe + H 2 SO 4  → FeSO 4  + H 2  ↑ (dilute) Copper lies below hydrogen . Thus, it cannot displace hydrogen from acids. Due to violent reactivity of K, Na, they are restricted in use with acids that causes explosion

21. Don’t write or place any image in this area ELECTROCHEMISTRY Nernst Equation Nernst Equation allows one to predict cell voltages (emf) when the conditions are not standard   R is universal gas constant (8.314 J K-1 mol-1 ) F is Faraday (96500 C) = Charge Carried by 1 mole of electrons T is temperature in Kelvin scale. n is number of electrons involved in electrode process. Activity is effective concentrations (in dilute solutions activity is molar concentration) The voltage difference between electrodes of an electrochemical cell is called cell voltage (or emf) E cell or E At standard conditions, E cell or E For a spontaneous reaction both E and E should be positive

22. Don’t write or place any image in this area ELECTROCHEMISTRY Derivation of Nernst Equation E red  = E Mn +/M = E o Mn +/M – [2.303RT/ nF ] log [1/[M n+ ]] General reaction M n+  + ne –  →  nM The calculation of single electrode reduction potential ( E red ) from the standard single electrode reduction potential ( E° red ) for an atom/ion is given by the Nernst equation. In a reversible reaction electrical energy is produced at expense of decrease in free energy Δ G = - nFE (1) Under standard conditions, Δ G = -nFE (2) The decrease in free energy Δ G associated with an equilibrium reaction is given by, Δ G = Δ G +RT ln Q (3) Substituting (1) and (2) in (3) - nFE = -nFE + RT ln Q On rearranging it, E = E ln Q  

23. Don’t write or place any image in this area ELECTROCHEMISTRY Nernst Equation Applications To find the concentration of one of the ionic species in cell if the concentration of the other ionic species is known It can be used to find the number of electrons involved in the redox reaction pH of the solution can be obtained pH = -log [H + ] To find the equilibrium constant of a reaction To study the e ffect of electrolyte concentration of single electrode potential E cell  = E  – [RT/ nF ] ln Q

24. Don’t write or place any image in this area ELECTROCHEMISTRY