POTENTIOMETRY.pdf
- 2. Introduction: Electrochemistry Electrochemistry is the study of the relationship between chemical transformation and electrical energy. Electrochemical cell It is the cell in which conversion of chemical energy to electrical energy can occur in either direction.
- 3. Electrochemical Cell Electrolytic Cell Galvanic or Voltaic Cell Converts electrical energy into chemical energy with non-spontaneous redox reaction. Converts chemical energy into electrical energy with spontaneous redox reaction. Electrochemical cell is of two types-
- 4. Electrochemical cell consists of two electrodes viz. anode (oxidation reaction) and cathode (reduction reaction) The electrode at which an oxidation reaction takes place is called as anode whereas the electrode at which a reduction reaction is takes place is called as cathode. Electrode Charges Cell Anode Cathode Electrolytic Cell Positive Negative Galvanic or Voltaic Cell Negative Positive
- 5. Ag Ag + + e- H+ + e- 1 2 H2 Redox Reaction: Oxidation: Loss of electrons causes increase in oxidation state. Reduction: Gain of electrons causes decrease in oxidation state.
- 6. Difference between Electrolytic and Galvanic or Voltaic Cells
- 7. Cell Notation of Electrochemical Cell: For drawing a cell diagram, we follow the following conventions. i. The anode is always placed on the left side, and the cathode is placed on the right side. ii. The salt bridge is represented by double vertical lines (||). iii. The difference in the phase of an element is represented by a single vertical line (|), while changes in oxidation states are represented by commas (,). iv. Concentration of aqueous solutions are written in parentheses after the symbol for ion or molecule.
- 8. Consider the following reaction: 2Ag+ (aq)+Cu(s) Cu2+ (aq)+2Ag(s) Two half Reactions At anode (Oxidation): Cu Cu+2 + 2e- At cathode (Reduction): Ag+ + e- Ag Half-Cell: When electrodes are immersed in a solution containing ions of the same metal, it is called a half-cell.
- 9. Salt Bridge: U-shaped tube containing an electrolyte (typically in the form of a gel), providing electrical contact between two solutions. The electrolyte used to prepare salt bridge is generally KNO3, KCl or NH4NO3
- 10. Cell Potential: (Ecell) Cell potential is the potential difference between two electrodes. Ecell = Ecathode – Eanode In Standard state, Eo cell = Eo cathode – Eo anode
- 11. Reference Electrode: A reference electrode is that electrode whose potential is known and remain constant. e.g. Saturated calomel electrode (ESCE = 0.242) Indicator Electrode: An indicator electrode is that electrode whose potential depends on the activity of ions being titrated or estimated. e.g. To carry out acid-base potentiometric titration Hydrogen gas. Quinhydrone electrode and glass electrodes are used as indicator electrode.
- 12. REFERENCE ELECTRODES • 1. Standard Hydrogen Electrode • 2. Saturated Calomel Reference Electrode • 3. Silver/Silver Chloride -Reference Electrode
- 13. Indicator Electrodes I. Metal Indicator Electrodes A. Electrodes of the First Kind B. Electrodes of the Second Kind C. Inert Metallic Electrodes (for Redox Systems) II. Membrane IE/Ion –Selective Electrodes. A. Glass pH IE B. Glass IE for other cations C. Liquid Membrane IE D. Crystalline-Membrane I III. Gas Sensing Probes/Gas Sensing Electrode.
- 14. Standard Hydrogen Electrode • A Standard Hydrogen Electrode (SHE) is referred as a primary standard electrode. • It can be consist of platinum foil coated with platinum black and has a wire contacts through a mercury. • This assembly is enclosed in a glass covering. o Hydrogen gas is passed at 1atmospheric pressure continuously. • The cell is represented as: Pt(s) H2(g), 1 atm H+ (aq.), 1M
- 15. • The cell potential for the standard hydrogen electrode is defined to be exactly zero. Advantages ➢It can be used over the entire pH. ➢It can be used as a reference electrode when dipped with standard acid solution and as a indicator electrode when dipped into a sample solution. ➢It is the primary reference standard against which the potentials of other electrodes are measured. Disadvantages ➢The Standard Hydrogen electrode is difficult to set up. ➢It is also difficult to handle. ➢It cannot be used in the solution containing oxidising and reducing agents. ➢It cannot be used also with unsaturated organic compounds and compounds containing sulphur and arsenic.
- 16. Saturated Calomel Electrode ➢It consist of inner jacket and outer sleeve ➢Inner tube has a wire contact with mercury(Hg) and plugged with a mixture of calomel(Hg2cl2)and KCL. ➢This is surrounded by an outer sleeve ➢Tip is filled with Crystals of KCL and porous plug of asbestos. ➢The space between inner jacket and outer sleeve is filled with saturated KCl/1N or KCl/0.1 N . ➢ Potential of the electrode depends on ➢The Concentration of the KCL Solution. And Temperature.
- 17. the half cell is represented as Hg Hg2Cl2 (satd), KCl (xM)|| The half cell reaction Hg2Cl2 (s) + 2e - 2Hg(l) + 2Cl- (aq) Advantages: ➢ Easy of construction ➢Stability of Potential. Disadvantages ➢Dependent on temp ➢Toxic Applications: Used in pH Measurement
- 18. Silver/Silver Chloride Electrode ➢Most widely used reference electrodes ➢Ag electrode immersed in KCl solution saturated with AgCl ➢ Potential depends on the concentration of the KCL and temperature. AgAgCl (satd),KCl (satd) AgCl(s) + e-Ag(s)+Cl-(aq) E = 0.199 V
- 19. Advantages Ag/AgCl electrode can be used at temperatures > 60oC Disadvantage Ag reacts with more ions, - plugging of the junction between electrode (Ag) and analyte soln
- 20. Indicator Electrodes 1. Electrodes of First kind • They are composed of the metal rod immersed in its metal solution. • These electrodes respond to the ionic activity of the electrode. Ex: silver electrode dipped in the silver nitrate solution. copper electrode dipped in the copper sulphate solution M+n(aq) + ne- M(s) Advantages • Not very selective • Ag+ interferes with Cu+2 • May be pH dependent • Zn and Cd dissolve in acidic solutions • Easily oxidized (deaeration required) • Non-reproducible response
- 21. Electrodes of Second Kind • These are composed of the metal wires coated with the salt precipitates. • These electrodes respond to the changes in the ionic activity through the formation of the complex or precipitates • Ex:1. Ag electrode for Cl- determination 2. Hg electrode for EDTA determination
- 22. • Electrodes of third Kind-Inert metallic Indicators • These electrodes are also known as inert electrodes and redox electrodes. • They are composed of inert metal electrode immersed in the solution containing the soluble oxidized and reduced forms of the redox half- reaction. • May not be reversible • Electrode acts as e- source/sink for electrons transferred from a redox system in the solution • Examples: Detection of Ce3+ with Pt electrode ½ reaction: Ce4+ + e - Ce3+
- 23. 2. Glass electrode: • The glass electrode belongs to category of ion selective electrodes. • It is characterized by presence of a glass membrane which shows selectivity for a particular ion. • The glass electrode used for the determination of pH shows a preferential response for hydrogen ions.
- 24. • The potential developed by the electrode is proportional to the difference in the hydrogen ion concentration on the both side of the glass membrane. • When the hydrogen ion concentration, on one side of the glass membrane is held constant, then the potential developed by the electrode will become proportional to the hydrogen ion concentration of the test solution. • The Ag(s)-AgCl(s) electrode provides the constant potential and the solution of 0.1N HCl provides the constant hydrogen ion concentration on one side of the glass membrane.
- 25. Construction: • The glass electrode consists of a glass tube ended into a glass bulb containing a solution of constant pH and electrode of constant potential (Ag-AgCl). • Special electronic circuits have to be devised for the measurement of the pH of the solution. • Instruments designed for this purpose, that make use of a glass and calomel electrode assembly and provide directly the pH of the solution are known as pH meter.
- 26. Mechanism: • This electrode involves the exchange of hydrogen ions of solution with silver ion of the glass electrode. • The extent of the exchange is depending on the concentration of H+ ion concentration in the respective solutions. • Due to this there is development of potential that is similar to the junction potential at each surface. • As long as the activity of water on the both sides of the glass membrane are equal, the response of the electrode remains proper. • Whenever ions other than hydrogen ion participate in the exchange process, the response of the electrode changes.
- 27. Electrode Potential: EG = E0 Glass - 0.05916 pH Cell Set-up: Glass Electrode Salt Bridge Calomel Electrode
- 28. Merits and Demerits of the electrode: Merits: i. It provides a measure of pH in the pH range of 1 – 9. ii. Using a pH meter, pH of the solution can be directly read. iii. The electrode can be used in all aqueous solutions. iv. Electrode is not affected by oxidizing and reducing agents or by any organic compound. v. pH can be determined even for small volume of solution.
- 29. Demerits: i. The electrode cannot function in highly acidic or alkaline medium ii. It cannot produce proper response with pH > 9 or <0.5. iii. It cannot function in non-aqueous medium. iv. It needs standardisation every time before the use.
- 30. Potentiometric Titrations: Principle: • Determination of equivalence point of titration on the basis of potential measurement using a suitable set up of galvanic cell is called as potentiometric titration. • Galvanic cell consists of two half cells. One half cell is called as indicator half cell and second half is called reference half cell. • During the potentiometric titration, the potential of indicator half cell changes whereas the potential of the reference half cell remains constant.
- 31. • The electrode potential on the activity or concentration of the ion with which the electrode is reversible is given by the Nernst equation. E = Eo - 2.303 RT nF log10 𝑎(𝑅𝑒𝑑𝑢𝑐𝑒𝑑 𝑆𝑡𝑎𝑡𝑒) 𝑎 (𝑂𝑥𝑖𝑑𝑖𝑠𝑒𝑑 𝑆𝑡𝑎𝑡𝑒) Where, E = Electrode potential of the indicator electrode Eo = Standard electrode potential of the indicator electrode n = Number of electrons involved in electrode reaction R = Gas Constant F = Faraday Constant (96500 Coulomb) a(Reduced State) = Product of activities of species involved in reduced state a(Oxidation State) = Product of activities of species involved in reduced state
- 32. During the potentiometric titration it is observed that, • Initially the emf changes gradually • Towards the equivalence point it is quite rapidly • After the equivalence point it is changes gradually The equivalence point can therefore determined by finding the quantity of titrant added at the point at which the rate of change of emf or potential is maximum.
- 33. The emf of the cell after each addition of titrant is measured. The graph of E vs V or ΔE/ΔV vs V or ΔE2/ΔV2 vs V is plotted from which equivalence point can be determined. ------------------------------------ EMF Volume of Titrant in cm3 ΔE/ΔV Volume of Titrant in cm3 Δ 2 E/Δ 2 V Volume of Titrant in cm3
- 34. Experimental Set Up: Diagram: • It consists of pair of electrodes. • One is acting as indicator electrode reversible to ion which is being estimated and second is reference electrode like SCE whose potential is known and remain constant. • Both the electrodes are dipped in the experimental solution. • A stirrer is added to stir the solution after each addition of titrant. • Electrodes are connected to the appropriate terminals of potentiometer. Constructions:
- 35. Procedure: • A known volume of the sample solution containing the ions to be titrated is placed in a beaker and indicator electrode reversible to the ions to be titrated is dipped in it along with SCE. • The electrodes then connected to the appropriate terminals of the potentiometer. • The emf of the cell thus set up is noted before adding any amount of the titrant. This Ecell corresponds to the zero volume of the titrant added. • The titrant is added from the burette in small increments and after addition of each increments the solution is stirred and emf is noted. • Initially, large increments of the titrant (1 to 2 cm3) may be added, but near the equivalence point where potential changes suddenly small increments of the titrant (0.1 to 0.2 cm3) should be added at a time.
- 36. • Since the potential changes are sharp near to the equivalence point, it is necessary to take large number of observations with small increments of the titrant before and after the equivalence point. • The equivalence point of the titration can be obtained by plotting any one of the graphs viz. E vs V or ΔE/ΔV vs V or ΔE2/ΔV2 vs V. ------------------------------------ EMF Volume of Titrant in cm3 ΔE/ΔV Volume of Titrant in cm3 Δ 2 E/Δ 2 V Volume of Titrant in cm3
- 37. Graphical Methods of determining the equivalence point in potentiometric titration of strong acid and strong base: Consider the titration of HCl (Strong acid) against NaOH (Strong base). In potentiometric titration, emf is measured for different volumes of titrant added. The equivalance point then can be determined graphically by any one of the following method 1. From the graph of Ecell vs Volume of the titrant (NaOH): Normal Curve Method EMF Volume of Titrant in cm3 • In this method the emf of the cell is plotted against the volume of titrant added. • The graph of Ecell vs volume of the titrant is always S shaped or sigmoid curve. • If the curve is vertical near the equivalence point, the equivalence point can be determined by bisecting the straight portion of the curve.
- 38. • When the curve is symmetrical, the midpoint of the steep portion of the curve is the equivalence point. • However, if the curve is not symmetrical, then the above method will not give correct results. In this case a method of parallel tangents or the circle fitting method is adopted to determine equivalence point. 2. From the graph of 𝚫𝐄 𝚫𝐕 vs Volume of the titrant (NaOH): First Derivative Curve ΔE/ΔV Volume of Titrant in cm3 Plot of 𝚫𝐄 𝚫𝐕 vs volume of titrant added shows two portions the ascending portion representing the increase in 𝚫𝐄 𝚫𝐕 and descending portion representing decrease in 𝚫𝐄 𝚫𝐕 . These two portions are then extra plotted which will intersect at the corresponding to the maximum value of 𝚫𝐄 𝚫𝐕 and this will be the equivalence point.
- 39. Δ 2 E/Δ 2 V 3. From the graph of 𝚫𝟐 𝐄 𝚫𝟐 𝐕 vs Volume of the titrant (NaOH): Second Derivative Curve ------------------------------------ Volume of Titrant in cm3 • This method is very accurate method for locating equivalence point of titration. • The method is based on the fact that the point at which the first derivatives is maximum. • At the same point, the second derivative is zero. Therefore, at the equivalence point 𝚫𝟐 𝐄 𝚫𝟐 𝐕 = 0. • The graph consists of two sets of 𝚫𝟐 𝐄 𝚫𝟐 𝐕 values. • One with positive values of 𝚫𝟐 𝐄 𝚫𝟐 𝐕 and second is with negative values of 𝚫𝟐 𝐄 𝚫𝟐 𝐕 . • The maximum positive values of 𝚫𝟐 𝐄 𝚫𝟐 𝐕 is joined with maximum negative values of 𝚫𝟐 𝐄 𝚫𝟐 𝐕 . • The line cut the volume axis at which 𝚫𝟐 𝐄 𝚫𝟐 𝐕 is zero which gives Veq.
- 40. Advantages of Potentiometric Titrations: 1. All the categories of acid-base titrations can be carried out by potentiometrically. 2. Acid-base titration in non-aqueous medium can also be carried out. 3. Precipitation titration can be carried out by using metal-metal ion electrode or metal-metal insoluble salt type electrode. 4. Complexometric titration can be carried out by using mercury as indicator electrode. 5. Many redox system can be carried out potentiometrically. Disadvantages of Potentiometric Titrations: 1. This method is time consuming. 2. It requires a large number of observations. 3. In this titration end point is obtained graphically.
- 41. Applications of Potentiometric Titrations: 1. To determine alkalinity and carbonate content in sea water 2. To determine the phosphoric acid content in aerated drinks. 3. To determine acetic acid content in commercial vinegar sample. 4. To determine dissociation constant of dibasic and tribasic acid. 5. To estimate small charges of CO2 or O2 content in natural waters.
- 42. • Potentiometric Titration are carried out for 1. Neutralization Reaction: indicator electrode-Glass Electrode Reference Electrodes-SCE or Ag-AgCl electrode Potential at equivalence point is given by E=k - 0.0592pH Examples-Dibasic and tribasic acids are titrated with alkali Titration of mixture of acids (CH3COOH+HCl)Vs NaOH
- 43. 2.Redox Titration Indicator Electrode is platinum wire or foil Reference Electrode is SCE or Ag-AgCl Potential of indicator electrode is function of ratio of concentration of oxidized and reduced forms of ions. Examples-Ferrous ammonium sulphate in dil H2SO4 Vs KMnO4 Sodium Arsenite Vs KBrO3 E = Eo - 2.303 RT nF log10 𝑎(𝑅𝑒𝑑𝑢𝑐𝑒𝑑 𝑆𝑡𝑎𝑡𝑒) 𝑎 (𝑂𝑥𝑖𝑑𝑖𝑠𝑒𝑑 𝑆𝑡𝑎𝑡𝑒)