Unit VI: Instrumental Methods of Analysis and Agricultural Technology PDF

Unit VI : Instrumental Methods of Analysis and Agricultural Technology DR. NIDHI SHARMA 2 COURSE OUTCOME (C 203.6) Examine the analyte using electro analytical techniques. 3  Programme Outcomes (POs) Engin...

Unit VI : Instrumental Methods of Analysis and Agricultural Technology DR. NIDHI SHARMA 2 COURSE OUTCOME (C 203.6) Examine the analyte using electro analytical techniques. 3  Programme Outcomes (POs) Engineering Graduates will be able to: PO 1: Engineering knowledge PO 2: Problem analysis PO 3: Design/development of solutions PO 4: Conduct investigations of complex problems: PO 5: Modern tool usage PO 6: The engineer and society PO 7: Environment and sustainability PO 8: Ethics PO 9: Individual and teamwork PO 10: Communication PO 11: Project management and finance PO 12: Life-long learning 4  Introduction to instrumental methods of analysis [A] Conductometry: Introduction, conductivity cell, conductivity measurement of soil, conductometric titrations of acid versus base with titration curve. [B] pHmetry: Introduction, Reference electrode, calomel electrode, indicator electrode, standardization of pH meter, pH measurement of soil pH metric titration of strong acid versus strong base with titration curve 5  A] Conductometry: Introduction, conductivity cell, conductivity measurement of soil, conductometric titrations of acid versus base with titration curve. 6  Conductometry Conductometry deals with the measurement of electric conductance of electrolyte solution such as acids, bases and salts. These solution conduct due to movement of ions towards opposite charges electrodes. Conductometry is used in neutralization titration, precipitation tit. , chemical kinetics and plant laboratories. Introduction Ohm’s law. Conductometric measurements. Factor affecting conductivity. Application of conductometry. Conductometric titration-: Introduction. Types of conductometric tiration. Advantages of conductometric titration. 7  1. Ohm’s law Ohm’s law is obeyed by Metallic conductors and solutions of electrolytes. Ohm’s law states that, the current (I) flowing through a conductor is directly proportional to the potential difference (V) applied across the conductor and is inversely proportional to the resistance (R ) of the conductor. 𝑉 𝐼= 𝑅 I - Current in Ampere V – Potential difference in volts R – Resistance in ohm  2. Conductance (C) – is the ease with which current flows through the conductor or solution is reciprocal of resistance C=1/R C - ohm-1 or mho 3. specific resistance (ρ) Resistance of any uniform conductor is R is directly proportional to length l and inversely proportional to area of cross section A. R α l/A ρ is specific resistance, if l = 1cm , and A = 1cm2 then R = ρ Therefore ρ = R *A / l Unit =ohm.cm ρ= specific resistance l is length cm A is area cm2  4. specific conductance Specific conductance= Cell constant x Obs. Conductance 1S = 1 ohm-1 = 1 mho = Ω-1 1 Sm-1= 1 ohm-1 m-1 = 1Ω-1m-1 Cell Constant The ratio of (l/a) in the equation is known as cell constant. It is constant for a given cell. Cell constant = l/a Unit = cm -1 Sp. Conductance = Measured conductance (obs) * cell constant  6. Equivalent conductance (λ) ❖To understand the meaning of equivalent conductance, imagine a rectangular trough with two opposite sides made of metallic conductor (acting as electrodes) exactly 1 cm apart, If 1 cm3 (1 mL) solution containing 1 gram equivalent of an electrolyte is places in this container is measured. λ= k V V is ml of solution containing one gm equivalent of electrolyte If the concentration of the solution is c g equivalent per liter, then the volume containing 1 g equivalent of the electrolyte will be 1000/C. So equivalent conductance, 𝜿 𝒙 𝟏𝟎𝟎𝟎 (𝒐𝒉𝒎−𝟏 𝒄𝒎−𝟏 ) 𝒄𝒎𝟑 𝝀= 𝜿𝑽 = = 𝑪 𝒈𝒎 𝒆𝒒𝒖𝒊𝒗. The unit of equivalent conductance, λ is ohm-1 cm 2 gm equi-1. 7. Molar conductance (µ)  MEASUREMENT OF ELECTROLYTIC CONDUCTANCE Let Rx be an unknown resistor, R1 and R2 two standard resistors, R3 an adjustable resistor and G a galvanometer. The bridge is connected to a source of power S, a battery, and a tapping key K is placed in the path to control the connections. In the bridge the total current is divided into two paths: i1 through R1 and R3, and i2 through R2 and Rx. Under the balancing conditions, the potential at points B and D must be the same, i.e. the ohmic voltage drop through the resistors R1 and R2 must be the same. Hence, the potential at B (GB) must be equal to potential at D (GD). GB = GD or i1 R1 = i2R2 Similarly, i1 R3 = i2 Rx  Conductivity Cell: Measurement of Conductance Measurement of conductance of a solution is done by using conductivity cell with alternating current. - The electrodes fitted in the cell are made up of platinum plates coated with platinum black. These plates are welded to platinum wires fused in two thin glass tubes. The contact with copper wires of the circuit is made by dipping them in mercury contained in the tubes. 15  The commonly used arrangement for the measurement of resistance of the conductance cell is shown in Fig. Instead of galvanometer, a head phone is used. 'AB' is manganese (Mn) wire slightly stretched over a meter rule graduated in milimeters. A sliding contact 'H' moves along this wire 'R' is a resistance 50x, 'C' is the conductivity cell containing electrolytic solution. 'I' is the induction coil from which alternating current is led When current is flowing, any unplugged in the resistance box R. contact H is moved until the sound phone is minimum. When this occurs, we have 16 Factors affecting Conductance 1. Concentration – Concentration increase conductance decreases. 2. Size of ions – smaller the size of ion greater the mobility and conductivity. eg H+, Na+, Li+ 3. Charge on the ions – Greater the charge greater the conductivity Na+, Mg+2, Al+3 4. Temperature – Higher the temperature higher the conductivity. Because at higher temperature kinetic energy of ions are more. A. Acid base conductometric titration 1) Strong Acid with Strong Base – HCl with NaOH (In burette NaOH and in conical flask HCl) : HCl being a strong acid dissociates completely to produce the highly mobile H+ and Cl- (Chloride ions). Hence, initially the conductivity is high. During titration, when NaOH is added from the burette to HCl, the highly mobile H+ are replaced by the less mobile Na+. Hence, the conductivity goes on decreasing till the end-point. At the end-point i.e. On complete neutralization the conductivity is due to Na+ and Cl ̄. If further NaOH is added after the end-point, the conductivity will increase due to the presence of Na+ Mobility of ion and OH- in the solution. H+ > Cl- H+ Cl- + Na+ OH- Na+ Cl- + H2O Na+ and Cl- minimum conductance 1 ml 1N NaOH = 36.5 mg of HCl OH- < H+  2) Weak Acid with Strong Base CH3COOH (in titration flask) with NaOH In the beginning Initially, the conductance of the solution (in burette) : is low because of the poor dissociation of the weak acid i.e. acetic acid. CH3COOH CH3COO¯ + H+ On addition of NaOH from the burette the conductance [CH3COOH] + [Na+ + OH-] CH3COO- + Na+ + H2O increases gradually due to increasing amount of salt (sodium acetate) which dissociates completely. The conductance gradually increases till the end-point. At the end-point the conductance is due to Na+ and CH3COO- (acetate ions). After the end-point on further addition of NaOH conductance increases rapidly due to fast moving OH- and Conductance Na+. The point of intersection of the two lines gives the end-point Calculation 1000 ml 1N NaOH = 60 gm of CH3COOH 1 ml 1N NaOH = 60 mg of CH3COOH  3. Weak base with Strong acid Consider the titration of HCl with NH4OH known volume of HCl is taken in the Titration flask and the burette is filled with NH4OH H+Cl- + NH4OH NH4+Cl- + H2O Initially, the conductance is high as HCl completely dissociates to produce the highly mobile H+ and Cl-. On addition of NH4OH, the fast-moving H+ are replaced by the slow moving NH4 + and hence the conductance decreases rapidly till the end-point. At the end-point the conductance is due to NH4 + and Cl-. After the end-point, on further addition of NH4OH there is not much change in the conductance as NH4OH dissociates to a small extent as NH4OH which is weak electrolyte. 1 ml 1N HCl = 35 mg of NH4OH 4. Weak acid with Weak base Consider the titration of acetic acid (taken in Titration flask) and ammonium hydroxide (filled in the burette). Initially, the conductance is low due to the poor dissociation of the weak acid (acetic acid). As the titration proceeds, the conductance starts increasing due to formation of the strong electrolyte ammonium acetate (CH3COO-NH4+ ). At the end-point the conductance is due to CH3COO- and NH4+. After the end-point the conductance remains almost constant because the excess base (NH4OH) added is a weak electrolyte. *The conductometric method is particularly suitable for such titrations which do not give a sharp end-point with indicators. Neutralization reaction CH3COOH + NH4OH (CH3COO¯+ NH4+) + H2O  [B] pHmetry: Introduction, Reference electrode, calomel electrode, indicator electrode, standardization of pH meter, pH measurement of soil pH metric titration of strong acid versus strong base with titration curve 22 Reference Electrode: Reference electrode is defined as the electrode which has stable and reproducible potential and completes the cell acting as half cell. Ideal reference electrode : a) is reversible and obeys the Nernst's equation b) exhibits a potential that is constant with time. c) returns to its original potential after being subjected to small currents. Purpose of reference electrode : 1. to complete the cell internal circuit 2. provide stable potential Types of Reference Electrodes 1. Primary Ref. Electrode - Standard Hydrogen Electrode 2. Secondary Ref. Electrode - Saturated Calomel Electrode, silver-silver chloride electrode, mercury-mercury sulphate electrode etc 1. Calomel (Hg2Cl2) Electrode It consists of an outer glass tube fitted with a frit at the bottom to permit electrical contact with solution. Inside this glass tube there is another tube having mercury at the bottom over which a paste of mercury-mercurous chloride (Hg - Hg2Cl2) is placed. A solution of potassium chloride (KCI) is then placed over the paste. A platinum wire sealed in glass tube helps in making the electrical contact. The electrode is connected with the help of the side tube on the left through a salt bridge with the other electrode to make a complete cell. Construction : or Hg/ Hg2Cl2. KCl (x M) If calomel acts as cathode Concentration E° Name of electrode of KCl calomel Hg2Cl2 + 2e¯ 2Hg + 2Cl¯ --------- (1) If calomel acts as anode 0.1 N 0.3334 Deci normal 2Hg +2Cl¯ Hg2Cl2 + 2e¯ -------- (2) volts calomel Ecal = Ecal° - 0.0591 log [ Cl¯] 1.0 N 0.2810 Normal calomel Demerits of Calomel Electrode volts 1. Not use above 80°C as Hg2Cl2 starts Saturated 0.2422 decomposing into Hg and HgCl2 Saturated calomel volts 2. Involves handling of poisonous Hg and Hg2Cl2 25  2. Indicator Electrode The electrode of a cell in which the potential depends on the concentration of particular ion is called as Indicator electrode. The potential of the indicator electrode changes depending on the concentration of the analyte. Example - Glass electrode Eg Composition of glass of Glass electrode – 72% SiO2, 21% Na2O, 6% CaO. Principle of Glass electrode : When two solution of different [H+] are separated by a thin glass membrane, a potential difference developed is proportional to the difference in [H+] of the two solutions. The glass electrode may be represented as, Ag, AgCl / 0.1 M HCl / H+ (unknown conc.) (test solution) Glass Electrode The glass electrode consists of a very thin walled glass bulb, made from a low melting point glass having high electrical conductivity, blown at the end of a glass tube as shown in the Figure. The bulb contains 0.1 M HCl solution, sealed into the glass- tube is a silver wire coated with silver chloride at its lower end. The lower end of the silver wire dipped into the HCl, silver-silver chloride electrode. When glass electrode is placed in a solution the potential develops across the glass membrane as a result of a concentration difference of H+ ions on the two sides of the membrane. The glass membrane acts as an ion exchanger i.e. exchange of Na+ of glass with H+ of solution. The potential of a glass electrode is determined using standard calomel electrode as shown in the Figure. 27  Determination of pH of Solution using Glass Electrode A glass electrode is couple with the calomel electrode to determine the pH of the solution Calomel H+ Glass electrode Unknown electrode Ecell= Ecal– EG = 0.2422 – (E°G+ 0.0591pH) pH =  Glass electrode require soaking of water for some hours before use so that – surface becomes active, hydrated layer developed on outer glass surface Ion exchange can take place easily. Advantages 1. It is stable electrode and can be used in presence of strong oxidizing and strong reducing agents. 2. It is compact and portable. 3. It attained equilibrium quickly. 4. It can detect and estimate H+ ion in presence of other ions. 5. simple to use  pH metric Measurement The term pH was introduced by Sorensen in 1909. pH = - log [H+] pOH = - log [OH-] pH is the negative logarthim of hydrogen ion concentration. [H+] = 10 – pH e.g. If pH of any acid solution is 3 then its [H+] will be 10-3 & [OH-] will be 10-11 Buffer Solution : A buffer solution maintains the pH fairly constant even upon the addition of small amount or acid or base. There are Two common types of Buffer solutions- 1. Acidic Buffer : These are prepared by using a Weak Acid with a salt of Weak Acid and Strong Base. e.g. CH3COOH + CH3COONa 2. Basic Buffer: These are prepared by using a Weak Base with a salt of Weak Base and Strong Acid. e.g. NH4OH + NH4Cl  pH metric titration of strong acid Procedure: Part I Standardization of pH meter A digital pH meter is calibrated as Connect the glass-calomel combine electrode to the pH meter and pH meter to electricity. Clean the electrode with distilled water. Keep the temperature knob on temperature of solution. Keep the functional knob on pH. Immerse the electrode in pH 7 solution. If digital display is not showing pH 7 then adjust by using calibration knob. Remove the electrode from pH 7 solution, clean with distilled water and deep in pH 4 solution. If digital display is not showing pH 4 then adjust by using slop knob. Remove the electrode and clean with distilled water. Repeat above steps to ensure the standardization  Part II Titration strong acid Vs Strong Base (HCl vs NaOH) Wash the electrode with distilled water. Pipette out 25 ml of HCl solution in the 100 ml beaker. Keep the magnetic stirrer on. Fill the burette with standard (0.1N) NaOH solution. Immerse the electrode in the solution in beaker. Note the pH at initial. Add 1 ml of NaOH to the acid in the beaker. Stir well and note down the Ph. Continue addition of 0.5 ml of NaOH and noting the pH till a sudden increase in pH is observed. Continue this till the pH value remains nearly constant. Reaction: NaOH+ HCl → NaCl + H2O Observation : From the graph, equivalence point at pH 7 for the titration is noted down as 'X' ml X ml = 0.1 N NaOH required for neutralization of HCl taken. Calculations : Using N1V1 (HCI) = N2 V2 (NaOH) concentration Normality of HCl is calculated  Strong acids vs Strong base Reactions 1) HCl +NaOH H2O+ NaCl Calculations: N1V1 = N2V2 HCl quantity HCl in 25 ml acid mixture = V1 ml of NaOH HCl in 1000 ml acid mixture = 1000/25 *V1 ml of NaOH= 40ml of NaOH If the normaity of NaOH is Z then, 1 ml 1N NaOH= 36.5 mg HCl per litre 1. pH value of water at 298 K is a) 0 b) 3 c) 7 d)10 -1 -1 -1 2. 1 Sm = ------ ohm m a) 1 b) 10 c) 0.1 d) 0.01 3. For saturated KCl solution, calomel electrode shows,--- a) E° calomel = 0.2422 V b) E° calomel = 0.2224 V c) E° calomel = 0.4242 V d) E° calomel = 0.2242 V 4. SI unit of specific conductance is --------------- -1 -1 2 -1 -1 -1 a) S m b) Sm c) S M d) S m 5. The unit conductance is -1 a) Ohm cm b) Ohm -1 c) Ohm cm d) Ohm 6. The ratio of sp. conductance to measured conductance is known as ….. a) Equivalent conductance b) molar conductance c) cell constant d) emf ❖ The pH of 0.001M HCl is ………. a) 0.001 b) -3 c) 11 d) 3 ❖ As concentration of solution decreases, its specific conductance --- a) Decreases b) increases c) remain constant d) none of these ❖ Acid base titration using pH metry glass electrode is coupled with ……… electrode a) Hydrogen c) Quinhydrone b) Calomel d) Ion selective ❖ If sp. Conductance and Observed conductance of a solution are same the cell Constant is equal to ………………..cm-1 a) Zero b) 0.5 c) 1.0 d) 10.0 ❖ If the pH of solution is 3, the [H+] is …………… moles/ litre a) 1 x 10-1 b) 2 x 10 -1 c) 1 x 10-3 d) 3 x 10-2 ❖ The factors that affect the mobility of ions in conductance measurement are a) Size b) charge c) Extent of solvation d) All of the above Thank You