Potentiometry.pptx

Potentiometry: Electrical potential, electrochemical cell, reference electrodes, indicator
electrodes, measurement of potential and Ph, construction and working of electrodes,
Potentiometric titrations, methods of detecting end point, Karl Fischer titration.

1. DISCOVER . LEARN . EMPOWER Potentiometry UNIVERSITY INSTITUTE OF PHARMA SCIENCES Pharm D (3 rd Year) Pharmaceutical Analysis Mr. Yunes Alsayadi Assistant Professor of Pharmaceutical Analysis

2. Potentiometry Potentiometry is one of the methods of electroanalytical chemistry. It is usually employed to find the  concentration  of a solute in solution. In potentiometric measurements, the potential between two electrodes is measured using a high impedance voltmeter.

3. History In potentiometry we measure the potential of an electrochemical cell under static conditions. Because no current—or only a negligible current—flows through the electrochemical cell, its composition remains unchanged. For this reason, potentiometry is a useful quantitative method of analysis. The first quantitative potentiometric applications appeared soon after the formulation, in 1889, of the Nernst equation, which relates an electrochemical cell’s potential to the concentration of electroactive species in the cell,

4. Potentiometry Principles The potential difference between the two electrodes used forms the basis of the potentiometry principle. The addition of a titrant leads to a change in the ionic concentration, causing changes in the potential difference.  The indicator electrode measures this potential difference. The reference electrode has a potential value and remains stable when dipped into the sample solution. The salt bridge is a divide used during potentiometric titration to avoid the intervention of the analyte solution with the reference solution. Analyte solution is the solution whose potential we must determine. 

5. We can calculate the total electric potential or the potential difference as: E cell  = E ind  – E ref  + E j Here,  E cell  = potential of the whole cell E ind  = potential of the indicator E ref  = potential or electromotive force of the reference electrode E j  = potential at the junction across the salt bridge The measured electrical potential depends on the concentration of the ions in contact with the indicator electrode.

6. Electrical Potential Potentiometry is based on the measurement of the potential of an electrode system (e.g. electrochemical cell). Potentiometric measurement system consists of two electrodes called reference and indicator electrode, potentiometer and a solution of analyte. Potential of an indicator electrode depends mainly on the concentration of the analyte ions. Potentiometric measurement system / Potentiometer (for pH measurement)

7. Electrochemical cell Electrochemical cell consists of two solutions connected by a salt bridge and electrodes to form electrical circuit. Sample cell consists of solutions of ZnSO4 and CuSO4. Metallic Zn and Cu electrodes are immersed in respective solutions. Electrodes have contacts firstly through wires connected to the voltmeter and secondly through solutions and a salt bridge, forming an electric circuit. Salt bridge consists of a tube filled with saturated salt solution (e.g. KCl solution). The ends of the tube are capped with porous frits that prevent solutions from mixing, but permit movement of ions. A galvanic electrochemical cell

8. Three distinct charge transfer processes are described for the system in Fig : 1. Electrons move in electrodes and wires from zinc electrode to copper electrode. 2. Ions move in solutions: a. In solution on the left, zinc ions move away from the electrode and sulfate ions move towards it. b. In solution on the right, copper ions move towards the electrode and negatively charged ions (sulfate) away from it. c. In salt bridge positive ions move right and negative ions left. 3. On the surfaces of electrodes electrons are transferred to ions or vice versa: a. Zinc electrode dissolves: Zn → Zn 2+ + 2e- b. Metallic copper is deposited on the electrode surface: Cu 2 + + 2e- → Cu ↓

9. Nernst equation Potential on an electrode depends on the ions present in the solution and their concentration. This way electrochemical cells can be used to determine ions and their concentration in solution. The dependence of potential between electrodes from concentration of ions is expressed by Nernst equation: E – electrode potential, E – standard potential of the electrode, R – universal gas constant (8.314 J/( K•mol )), F – Faraday constant (96485 C/mol), T – temperature in kelvins, n – charge of the ion or number of electrons participating in the reaction, a – activity of the ions.

10. Reference electrodes T he electrode whose half-cell potential is known and is constant and completely insensitive to the composition of the solution is called a reference electrode. The reference electrode can act as both anode or cathode depending upon the nature of other electrodes. A reference electrode has a stable and well defined electrochemical potential (at constant temperature), against which the applied or measured potentials in an electrochemical cell are referred. The reference electrodes are classified into two types: Primary reference electrode:  The standard hydrogen electrode is called a primary reference electrode. Secondary reference electrode:  The electrode whose potential is determined by connecting to the standard hydrogen electrode is called a secondary reference electrode. Example: Calomel electrode

11. Standard Hydrogen Electrode (SHE) Standard hydrogen electrode scheme: 1) Platinized platinum electrode, 2) Hydrogen gas, 3) Acid solution with an activity of H +  = 1 mol/L, 4) Hydroseal for prevention of oxygen interference, 5) Reservoir via which the second half-element of the galvanic cell should be attached. The connection can be direct, through a narrow tube to reduce mixing, or through a  salt bridge , depending on the other electrode and solution. This creates an ionically conductive path to the working electrode of interest.

12. Standard Hydrogen Electrode (SHE) It is rarely used for routine analytical work, but is important because it is the reference electrode used to establish standard‑state potentials for other half‑reactions. A conventional salt bridge connects the SHE to the indicator half‑cell. The shorthand notation for the standard hydrogen electrode is Pt(s), H 2 (g, 1 atm) II H + ( aq ) 2H + ( aq ) + 2e  H 2 (g); E o = 0 V

13. Saturated Calomel Electrode (SCE) The  saturated calomel electrode  (SCE) is a  reference electrode  based on the reaction between elemental  mercury  and  mercury(I) chloride . It has been widely replaced by the  silver chloride electrode , however the calomel electrode has a reputation of being more robust. The aqueous phase in contact with the mercury and the mercury(I) chloride (Hg 2 Cl 2 , " calomel ") is a saturated solution of  potassium  chloride in water. The electrode is normally linked via a porous  frit  to the solution in which the other electrode is immersed. This  porous  frit is a  salt bridge . The calomel electrode is very practical and very robust, and is one of the most common electrodes used in corrosion studies. The SCE is used in pH measurement, cyclic voltammetry and general aqueous electrochemistry. The SCE met the criteria of ideal reference electrodes; therefore, SCE is the most popular reference electrode for use in aqueous solutions.

14. Working The electrode is represented as : KCl ( satu .) Hg 2 Cl 2(s) Hg (l) Oxidation : If the electrode serves as anode, then half reaction that occurs on it will be oxidation. Mercury is first oxidised to mercuric ions. 2Hg (l) ⟶Hg 2 +2 +2e − The chloride ions supplied by  KCl  solution combine with mercuric ions [Hg 2 +2 ] to form insoluble mercurous chloride. Thus, Hg 2 2+ +2Cl − ⟶Hg 2 Cl 2(s) Overall reaction is, 2Hg (l) +2Cl − ( aq ) ⟶Hg 2 Cl 2(s) +2e − Reduction : If the electrode is cathode in the galvanic cell, the half reaction that occurs on it will be reduction : Hg 2 Cl 2(s) +2e−⟶2Hg (l) +2Cl − ( aq ) The potential of calomel electrode decreases with increase in the concentration of chloride ions at a given temperature. Thus, electrode is reversible with respect to concentration of chloride ions.

15. The shorthand notation for the calomel electrode half-cell is: Hg(l) | Hg 2 Cl 2 ( sat'd ), KCl ( aq , saturated) The SCE has the advantage that the concentration of Cl ‑ , and, therefore, the potential of the electrode, remains constant even if the KCl solution partially evaporates. A significant disadvantage of the SCE is that the solubility of KCl is sensitive to a change in temperature. At higher temperatures the concentration of Cl ‑ increases, and the electrode's potential decreases. Electrodes containing unsaturated solutions of KCl have potentials that are less temperature‑dependent, but experience a change in potential if the concentration of KCl increases due to evaporation. Another disadvantage to calomel electrodes is that they cannot be used at temperatures above 80 °C.

16. INDICATOR ELECTRODE The electrode is used to measure the potential of a solution to which it is dipped inside along with reference electrode. Indicator Electrode : electrode that responds to analyte and donates/accepts electrons while Reference Electrode : second ½ cell at a constant potential Cell voltage is difference between the indicator and reference electrode .

17. Indicator Electrodes 1.) Two Broad Classes of Indicator Electrodes Metal Electrodes Develop an electric potential in response to a redox reaction at the metal surface Ion-selective Electrodes Selectively bind one type of ion to a membrane to generate an electric potential Remember an electric potential is generated by a separation of charge

18. METALLIC INDICATOR ELECTRODES

19. Electrodes of the First Kind Pure metal electrode in direct equilibrium with its cation Metal is in contact with a solution containing its cation. M + n ( aq ) + n e -  M(s)

20. Disadvantages of First Kind Electrodes 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 the Second Kind Respond to anions by forming precipitates or stable complex Examples: Ag electrode for Cl - determination Hg electrode for EDTA determination

22. Inert Metallic (Redox) Electrodes Inert conductors that respond to redox systems Electron source or sink An inert metal in contact with a solution containing the soluble oxidized and reduced forms of the redox half-reaction. May not be reversible Examples: Pt, Au, Pd, C

23. MEMBRANE ELECTRODES Consist of a thin membrane separating 2 solutions of different ion concentrations Most common: pH Glass electrode

24. Glass pH Electrode

25. Properties of Glass pH electrode Potential not affected by the presence of oxidizing or reducing agents Operates over a wide pH range Fast response Functions well in physiological systems Very selective Long lifespan

26. Liquid Membrane Electrodes Potential develops across the interface between the analyte solution and a liquid ion exchanger (that bonds with analyte) Similar to a pH electrode except that the membrane is an organic polymer saturated with a liquid ion exchanger Used for polyvalent ions as well as some anions Example: Calcium dialkyl phosphate insoluble in water, but binds Ca 2+ strongly

27.                                                                              Responsive to Ca 2+ 0.1 M CaCl 2

29. Crystalline-Membrane Electrodes Solid state electrodes Usually ionic compound Crushed powder, melted and formed Sometimes doped to increase conductivity Operation similar to glass membrane

30. Crystalline-Membrane Electrodes AgX membrane: Determination of X - Ag 2 S membrane: Determination of S -2 LaF 3 membrane: Determination of F -

31. Gas Sensing Probes A galvanic cell whose potential is related to the concentration of a gas in solution Consist of RE, ISE and electrolyte solution A thin gas-permeable membrane (PTFE) serves as a barrier between internal and analyte solutions Allows small gas molecules to pass and dissolve into internal solution O 2 , NH 3 /NH 4 + , and CO 2 /HCO 3 - /CO 3 2-

32. Gas Sensing Probe

33. 34 Potentiometric Titration Method Potentiometric Titration is done via the usage of two  electrodes  – an indicator electrode and a reference electrode (generally a hydrogen electrode or a silver chloride electrode). One half-cell is formed with the indicator electrode and the ions of the analyte , which is generally an electrolyte solution. The other half-cell is formed by the reference electrode. The overall cell potential can be calculated using the formula given below. E cell = E ind – E ref +E sol Where the potential drop between the indicator and reference electrodes over the electrolyte solution is given by E sol . The overall cell potential, E cell  is calculated in every interval where the titrant is measured and added. Now, a graph is plotted with the Potential difference on the Y-axis and the volume on the X-axis as shown below. It can be observed from the graph that the electric potential of the cell is dependent on the concentration of ions which are in contact with the indicator electrode. Therefore, the E cell   is measured with each addition of the titrant.

34. 35 Types of Potentiometric Titration There are four types of titration that fall under the category of potentiometric titration, namely acid-base titration, redox titration, complexometric titration, and  precipitation titration . A brief description of each of these types of titration is given below. 1. Acid-Base Titration:  This type of potentiometric titration is used to determine the concentration of a given acid/base by neutralizing it exactly using a standard solution of base/acid whose concentration is known. 2. Redox Titration:  This type of potentiometric titration involves an analyte and titrant that undergo a  redox reaction . An example of this type of titration would be the treatment of an iodine solution with a reducing agent which produces iodide ion (a starch indicator is used to get the endpoint). 3. Complexometric Titration:  This type of titration can also be referred to as chelatometry . In this method, a coloured complex is formed, indicating the end point of the titration. This method is used to determine a mixture of metal ions in a given solution. 4. Precipitation Titration:  This type of titration involves a reaction between the given analyte and the titrant wherein an insoluble  precipitate  is formed. The end-point of this titration is noted when the addition of the titrant no longer forms a precipitate.

35. POTENTIOMETRIC TITRATION Involves measurement of the potential of a suitable indicator electrode as a function of titrant volume Provides MORE RELIABLE data than the usual titration method Useful with colored/turbid solutions May be automated More time consuming

36. End point The endpoint of the titration is the point at which the colour of the solution changes completely due to the formation of product due to the addition of indicator. It is to be noted that weak acids only show one endpoint in the titration. Example: In acid base titration, due to addition of Phenolphthalein, the solution changes its colour . The main difference between equivalence and endpoint is that the equivalence point is a point where the chemical reaction comes to an end, while the endpoint is the point where the colour change occurs in a system. since there is only a slight difference between an equivalent point and an endpoint, it can be considered the same for laboratory purposes.

37. Methods of detecting end point Visual Inspection: In some cases, the endpoint can be determined visually by observing a sudden change in the color of an indicator solution. For example, phenolphthalein changes from colorless to pink in acidic solutions, indicating the endpoint in acid-base titrations. pH Measurement: pH meters can be used to monitor the pH of the solution throughout the titration. A significant change in pH often corresponds to the endpoint, especially in acid-base titrations.

38. Considerations for Specific Titrations Different types of potentiometric titrations require specific considerations for endpoint determination: Acid-Base Titrations: In acid-base titrations, the endpoint is often detected using pH indicators or pH meters. The pH at the equivalence point is typically close to 7 for neutralization reactions. Redox Titrations: In redox titrations, the endpoint is determined by monitoring changes in the voltage due to the transfer of electrons between the analyte and titrant. Complexometric Titrations: In complexometric titrations, the endpoint is typically detected using indicators or electrodes sensitive to metal ions. Ethylenediaminetetraacetic Acid (EDTA) is the example

39. Karl Fischer titration Karl Fischer (March 24, 1901 – April 16, 1958) was a German chemist Published a method in 1935 to determine trace amounts of water in samples. This method is now called Karl Fischer titration. Abbreviations: KF or KFT It remains the primary method of water content determination used worldwide by: Government – Food Science Academia – Research – Industry – Quality Control Karl Fischer titration  is a classic  titration  method that uses  coulometric  or  volumetric  titration to determine trace amounts of  water  in a sample.

40. Chemical principle The principle of Karl Fischer titration is based on the oxidation reaction between iodine and sulphur dioxide. Water reacts with iodine and sulphur dioxide to form sulphur trioxide and hydrogen iodide. An endpoint is reached when all the water is consumed. The chemical equation for the reaction between sulphur dioxide, iodine, and water (which is employed during Karl Fischer titration) is provided below. I 2  + SO 2  + H 2 O → 2HI + SO 3 This elementary reaction consumes exactly one molar equivalent of water vs. iodine. Iodine is added to the solution until it is present in excess, marking the end point of the titration, which can be detected by potentiometry. The reaction is run in an alcohol solution containing a base, which consumes the  sulfur trioxide  and  hydroiodic acid  produced.

41. Karl Fischer Titration Equipment Drying tube, sample injection cap, electrode analysis, Drain cook, a cathode chamber, detection electrode, rotor, anode chamber, KF reagent. Ingredients of KF reagent: Iodine, Buffer (Imidazole), sulphur dioxide, solvent (methanol).

42. Karl Fischer Titration Procedure The Karl Fischer titration experiment can be performed in two different methods. They are: Volumetric determination – This technique is suitable to determine water content down to 1% of water. The sample is dissolved in KF  methanol  and the iodine is added to KF Reagent. The endpoint is detected potentiometrically. Coulometric determination – The endpoint is detected in this experiment electrochemically. Iodine required for KF reaction is obtained by anodic oxidation of iodide from solution.

43. Advantages of Karl Fischer Titration It is fitted for determining water in gases, liquids and solids. The coulometric titrator helps in detecting free water, dissolved water, and emulsified water. It is a swift process which demands a minimal amount of sample preparation. Extremely accurate method. Limitations of Karl Fischer Titration It is a destructive technique. The solvent consumption is high as the manual volumetric titration demands reloading during each determination. Coulometric titration is fitted only for samples that contain a small amount of water. Coulometric titration takes extremely long periods to determine.

44. A.I. Vogel, Text Book of Quantitative Inorganic Analysis; 5 th Edition, John Wiley & Sons, Inc., New York. A.H. Beckett and J.B. Stenlake's , Practical Pharmaceutical Chemistry Vol I and II, Stahlone Press of University of London. P. Gundu Rao , Inorganic Pharmaceutical Chemistry. Bentley and Driver's Textbook of Pharmaceutical Chemistry. http://www.anachem.umu.se/jumpstation.htm http://userwww.service.emory.edu/~kmurray/mslist.html http://www.anachem.umu.se/jumpstation.htm References

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