Redox titrations & potentiometry

Redox titrations & potentiometry
(Mark=3) 920316 slides

http:\asadipour.kmu.ac.ir 81 slides
Redox reactions 2Fe3+ + Sn2+  2Fe2+ + Sn4+ half-reactions: Reduction 2Fe3+ + 2e-  2Fe2+ oxidation Sn2+  Sn4+ + 2e- 920316 slides

Standard Reduction Potentials (strength)
Reduction Half-Reaction E(V) F2(g) + 2e-  2F-(aq) 2.87 Au3+(aq) + 3e-  Au(s) 1.50 Cl2(g) + 2 e-  2Cl-(aq) 1.36 Cr2O72-(aq) + 14H+(aq) + 6e-  2Cr3+(aq) + 7H2O 1.33 O2(g) + 4H+ + 4e-  2H2O(l) 1.23 Ag+(aq) + e-  Ag(s) 0.80 Fe3+(aq) + e-  Fe2+(aq) 0.77 Cu2+(aq) + 2e-  Cu(s) 0.34 Sn4+(aq) + 2e-  Sn2+(aq) 0.15 2H+(aq) + 2e-  H2(g) 0.00 Sn2+(aq) + 2e-  Sn(s) -0.14 Ni2+(aq) + 2e-  Ni(s) -0.23 Fe2+(aq) + 2e-  Fe(s) -0.44 Zn2+(aq) + 2e-  Zn(s) -0.76 Al3+(aq) + 3e-  Al(s) -1.66 Mg2+(aq) + 2e-  Mg(s) -2.37 Li+(aq) + e-  Li(s) -3.04 Ox. agent strength increases Red. agent strength increases 920316 slides 3

Balancing of redox reactions. Under Acidic conditions
1. Identify oxidized and reduced species Write the half reaction for each. 2. Balance the half rxn separately except H & O’s. Balance: Oxygen by H2O Balance: Hydrogen by H+ Balance: Charge by e - 3. Multiply each half reaction by a coefficient. There should be the same # of e- in both half-rxn. 4. Add the half-rxn together, the e - should cancel. 920316 slides

Balancing of redox reactions. Under Basic conditions
1. Identify oxidized and reduced species Write the half reaction for each. 2. Balance the half rxn separately except H & O’s. Balance: Oxygen by H2O Balance: Hydrogen by OH- Balance: Charge by e - 3. Multiply each half reaction by a coefficient. There should be the same # of e- in both half-rxn. 4. Add the half-rxn together, the e - should cancel. 920316 slides

Balancing of redox reactions
H2O2 (aq) + Cr2O7-2(aq )  Cr 3+ (aq) + O2 (g) Redox reaction ====================================== 1)write 2 half reactions Half Rxn (red): Cr2O7-2 (aq)  Cr3+ Half Rxn (oxid): H2O2 (aq)  O2 2)Atom balance Cr2O7-2 (aq)  2Cr3+ 920316 slides

3)Oxygen balance Half Rxn (red): Cr2O7-2 (aq)  2Cr H2O Half Rxn (oxi): H2O2 (aq)  O2 4)Hydrogen balance Half Rxn (red): 14H+ + Cr2O7-2 (aq)  2Cr H2O Half Rxn (oxi): H2O2 (aq)  O H+ 5)Electron balance 6e H+ + Cr2O7-2 (aq)  2Cr H2O H2O2 (aq)  O H e- 920316 slides

6) Equalize of produced and consumed electrons 6e H+ + Cr2O7-2 (aq)  2Cr H2O ( H2O2 (aq)  O H e- ) x 3 7)Multiply each half reaction 8 H H2O2 + Cr2O72-  2Cr O H2O 920316 slides

Balance the redox reactions
I2 +S2O32- ⇋ I- +S4O62- I2 +S2O32- ⇋ I- +SO42- H+ OH- 920316 slides

Preadjustment of analyte oxidation state It is necessary to adjust the oxidation state of the analyte to one that can be titrated with an auxiliary oxidizing or reducing agent. Ex Preadjustment by auxiliary reagent Fe(II), Fe(III) Fe(II) – 4 Titration Ce4+ Preoxidation : Peroxydisulfate ( (NH4)2S2O8 ) 2 – ) Sodium bismuthate ( NaBiO 3 Hydrogen peroxide (H2O2) Prereduction : Stannous chloride ( SnCl2) Chromous chloride Jones reductor (zinc coated with zinc amalgam) Walden reductor ( solid Ag and 1M HCl) 920316 slides

Redox titrations 2Fe3+ + Sn2+  2Fe2+ + Sn4+ 2) In electrochemical cell. (Potentiometry) In solution (visual indicators) 920316 slides

Oxidizing and reducing agents 920316 slides

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Iodimetry and iodometry
a reducing analyte is titrated directly with iodine (to produce I−). iodometry : an oxidizing analyte is added to excess I− to produce iodine, which is then titrated with standard thiosulfate solution. I- + Cu2+→ I2 + Cu+ I2 + S2O32- → 2I- + S4O62- 920316 slides

Preparation of aqoues solution I3-
1) Iodine only dissolves slightly in water. Its solubility is enhanced by interacting with I- Standardization of Iodin with Arsenious oxide, As2O H2O = 2H3AsO3 As4O H2O = 4H3AsO3 H3AsO3 + I3– + H2O = H3AsO4 + 3I– + 2H+ 920316 slides

standard I3- 1) An excellent way to prepare standard I3- is to add a weighed quantity of potassium iodate to a small excess of KI. Then add excess strong acid (giving pH ≈ 1) to produce I3- by quantitative reverse disproportionation: 2) Cu + HNO3  Cu2+ Cu2++4I- 2CUI + I2 920316 slides

Stability of I2 Solutions
In acidic solutions of I3- are unstable because the excess I− is slowly oxidized by air: In neutral solutions, oxidation is insignificant in the absence of heat, light, and metal ions. At pH ≳ 11, triiodide disproportionates to hypoiodous acid (HOI), iodate, and iodide. I2 + OH- ⇌ IO- + I- + H+ 3IO- ⇌ IO3- + 2I- 920316 slides

Iodimetry 920316 slides

iodometry 920316 slides

Sodium thiosulfate, Na2S2O3 Thiosulfate ion is a moderate reducing agent that has been widely used to determine oxidizing agents by an indirect procedure that involves iodine as an intermediate. With iodine, thiosulfate ion is oxidized quantitatively to tetrathionate ion according to the half-reaction: 2S2O3 2–  S4O6 2– + 2e Eo = 0.08 Ex. Determination of hypochlorite in bleaches [CaCl(OCl)H2O]: OCl– + 2I– + 2H+  Cl– + I2 + H2O (unmeasured excess KI) I S2O3 2–  2I– + S4O6 2– Indicator: soluble starch (-amylose) 920316 slides

Standardization of thiosulfate solution: Primary standard : potassium iodate (KIO3), K2Cr2O7, KBrO3 Titration reactions: KIO3 + 5KI + 6HCl  3I KCl H2O I Na2S2O3  2NaI + Na2S4O6 KIO  3I  6Na2S2O3·5H2O  6 Equivalent S2O32- +H+ ⇋ HSO3- +S(s) pH, Microorganisms, Concentration, Cu2+, Sunlight Stabilizer for sodium thiosulfate solution : Na2CO3 Na2S2O3 + H2O + CO2  Na2CO3 + H2S2O3 H2S2O3  H2SO3 + S 920316 slides

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16-2 Finding the end point A redox indicator is a compound that changes color when it goes from its oxidized to its reduced state. or For ferroin, with E° = V we expect the color change to occur in the approximate range 1.088 V to V with respect SHE 920316 slides

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Starch-Iodine Complex Starch is the indicator of choice for those procedures involving iodine because it forms an intense blue colour with iodine. Starch is not a redox indicator; it responds specifically to the presence of I2, not to a change in redox potential. Structure of the repeating unit of the sugar amylose. 920316 slides

Permanganate titration Oxidation with permanganate : (Reduction of permanaganate) KMnO4 Powerful oxidant that the most widely used. 1) In strongly acidic solutions (1M H2SO4 or HCl, pH  1) MnO4– + 8H+ + 5e = Mn H2 O Eo = 1.51 V KMnO4 is a self-indicator. 2) In feebly acidic, neutral, or alkaline solutions MnO4– + 4H+ + 3e = MnO2 (s) + 2H2 O Eo = V 3) In very strongly alkaline solution (2M NaOH) MnO4– + e = MnO42 – Eo = V   920316 slides

Permanganate titration Duration of colour in end point (30 seconds) MnO4– + 3Mn2+ + 2H2O  5MnO2 + 4H K=1*1047 Stability of aqoues solution of MnO4- MnO4– + 2H2O  4MnO2 (s) + 3O2 (g) +4OH-   920316 slides

Standardization of KMnO4 solution Potassium permanganate is not primary standard, because traces of MnO2 are invariably present. Standardization by titration of sodium oxalate (primary standard) : 2KMnO Na2(COO)2 + 8H2SO4 = 2MnSO4 + K2SO4 + 5Na2SO CO2 + 8H2O 2KMnO  Na2(COO)  10 Equivalent 920316 slides

Preparation of 0.1 N potassium permanganate solution KMnO4 is not pure. Distilled water contains traces of organic reducing substances which react slowly with permanganate to form hydrous managnese dioxide. Manganesse dioxide promotes the autodecomposition of permanganate. 1) Dissolve about 3.2 g of KMnO4 (mw=158.04) in 1000ml of water, heat the solution to boiling, and keep slightly below the boiling point for 1 hr. Alternatively , allow the solution to stand at room temperature for 2 or 3 days. Filter the liquid through a sintered-glass filter crucible to remove solid MnO2. Transfer the filtrate to a clean stoppered bottle freed from grease with cleaning mixture and standardize it. Protect the solution from evaporation, dust, and reducing vapors, and keep it in the dark or in diffuse light. If in time managanese dioxide settles out, refilter the solution and restandardize it. 920316 slides

Applications of permanganometry H2O2 2KMnO H2O2 + 3H2SO4 = 2MnSO4 + K2SO4 + 5O2 + 8H2O (2) NaNO2 2NaNO2 + H2SO4 = Na2SO HNO2 2KMnO HNO2 + 3H2SO4 = 2MnSO4 + K2SO4 + 5HNO3 + 3H2O (3) FeSO4 2KMnO FeSO4 + 8H2SO4 = 2MnSO4 + K2SO4 + 5Fe2(SO4)3 + 8H2O (4) CaO CaO HCl = CaCl2 + H2O CaCl H2C2O4 = CaC2O HCl (excess oxalic acid) 2KMnO H2C2O4 + 3H2SO4 = 2MnSO4 + K2SO4 + 10CO2 + 8H2O (back tit) (5) Calcium gluconate [CH2OH(CHOH)4COO]2Ca HCl = CaCl CH2OH9CHOH)4COOH (NH4)2C2O CaCl2 = CaC2O NH4Cl CaCl H2SO4 = H2C2O4 + CaSO4 2KMnO H2C2O4 + 3H2SO4 = 2MnSO4 + K2SO4 + 10CO2 + 8H2O 920316 slides

Bromatimetry BrO3– + 5Br– + 6H+  3Br2 + H2O 2I– + Br2  I2 + 2Br– I S2O32–  2I– + S4O62– 920316 slides

Determining water with the Karl Fisher Reagent The Karl Fisher reaction : I SO H2O  2HI + H2SO4 For the determination of small amount of water, Karl Fischer(1935) proposed a reagent prepared as an anhydrous methanolic solution containing iodine, sulfur dioxide and anhydrous pyridine in the mole ratio 1:3:10. The reaction with water involves the following reactions : C5H5N•I2 + C5H5N•SO2 + C5H5N + H2O  2 C5H5N•HI + C5H5N•SO3 C5H5N+•SO3– + CH3OH  C5H5N(H)SO4CH3 Pyridinium sulfite can also consume water. C5H5N+•SO3– + H2O  C5H5NH+SO4H– It is always advisable to use fresh reagent because of the presence of various side reactions involving iodine. The reagent is stored in a desiccant-protected container. The end point can be detected either by visual( at the end point, the color changes from dark brown to yellow) or electrometric, or photometric (absorbance at 700nm) titration methods. The detection of water by the coulometric technique with Karl Fischer reagent is popular. 920316 slides

Potentiometric Methods A.) Introduction: 1.) Potentiometric Methods: based on measurements of the potential of electrochemical cells in the absence of appreciable currents (I →0) 2.) Basic Components: a) reference electrode: gives reference for potential measurement b) indicator electrode: where species of interest is measured c) potential measuring device 920316 slides

Electrodes and Potentiometry
Potential change only dependent on one ½ cell concentrations Reference electrode is fixed or saturated  doesn’t change! Ecell=Ecathod-Eanod Anod is conventionally reference electode Fe3+ +e- Fe2+ AgCl(s) + e- → Ag + Cl- Reference electrode, [Cl-] is constant Potential of the cell only depends on [Fe2+] & [Fe3+] Unknown solution of [Fe2+] & [Fe3+] Pt wire is indicator electrode whose potential responds to [Fe2+]/[Fe3+] 920316 slides

Reference Electrodes: (Instead of SHE) Need one electrode of system to act as a reference against which potential measurements can be made  relative comparison. Standard hydrogen electrodes are cumbersome Requires H2 gas and freshly prepared Pt surface Desired Characteristics: a) known or fixed potential b) constant response c) insensitive to composition of solution under study d) obeys Nernest Equation e) reversible > > > 920316 slides

Reference Electrodes 1.) Silver-Silver Chloride Reference Electrode Eo = V Activity of Cl- not 1E(sat,KCl) = V 920316 slides

Reference Electrodes 2.) Saturated Calomel Reference Electrode (S.C.E) Saturated KCl maintains constant [Cl-] even with some evaporation Eo = V Activity of Cl- not 1E(sat,KCl) = V 920316 slides

Indicator Electrodes 2 Broad Classes of Indicator Electrodes 1) Metal indicator Electrodes Develop an electric potential in response to a redox reaction at the metal surface 2) Membrane Indicator Electrodes a) Ion-selective Electrodes Selectively bind one type of ion to a membrane to generate an electric potential b) Molecular Selective Electrode 920316 slides

1) Metallic Indicator Electrode (3 Main Types) a) Metallic Electrodes of the First Kind b) Metallic Electrodes of the Second Kind c) Metallic Redox Indicators i. Involves single reaction ii. Detection of cathione derived from the metal used in the electrode iii. Example: use of copper electrode to detect Cu2+ in solution ½ reaction: Cu2+ + 2e-  Cu (s) Eind gives direct measure of Cu2+: Eind = EoCu – (0.0592/2) log aCu(s)/aCu2+ since aCu(s) = 1: Eind = EoCu – (0.0592/2) log 1/aCu2+ or using pCu = -log aCu2+: Eind = EoCu – (0.0592/2) pCu 920316 slides

b) Metallic Electrodes of the Second Kind i. Detection of anion derived from the interaction with metal ion (Mn+) from the electrode ii. Anion forms precipitate or stable complex with metal ion (Mn+) iii. Example: Detection of Cl- with Ag electrode ½ reaction: AgCl(s) + e-  Ag(s) + Cl- EO = V Eind gives direct measure of Cl-: Eind = Eo – (0.0592/1) log aAg(s) aCl-/aAgCl(s) since aAg(s) and aAgCl(s)= 1 & Eo = V: Eind = – (0.0592/1) log aCl- 920316 slides

c) Metallic Redox Indicators i. Electrodes made from inert metals (Pt, Au, Pd) ii. Used to detect oxidation/reduction in solution iii. Electrode acts as e- source/sink iv. Example: Detection of Ce3+ with Pt electrode ½ reaction: Ce4+ + e-  Ce3+ Eind responds to Ce4+: Eind = Eo – (0.0592/1) log aCe3+/aCe4+ 920316 slides

2) Membrane Indicator Electrodes i. electrodes based on determination of Molecules,cations or anions by the selective adsorption of these species to a membrane surface. ii. Ion Selective Electrodes (ISE) or pIon Electrodes are more common. iii. Desired properties of ISE’s 1) minimal solubility – membrane will not dissolve in solution during measurement. – silica, polymers, low solubility inorganic compounds , (AgX) can be used 2) Need some electrical conductivity 3) Selectively binds ion of interest 920316 slides

Ion-Selective Electrodes Does not involve a redox process Responds Selectively to one ion such as (C+) Contains a thin membrane capable of only binding the desired ion electric potential is generated by a separation of charge ISE + IRE = Combined electrod Constant C+ - Inner solution E2 Constant C+ - membrane E1 C+ - Un known solution E=E1-E2 920316 slides

Ion-Selective Electrodes Indicator Electrodes Ion-Selective Electrodes Responds Selectively to one ion Does not involve a redox process Contains a thin membrane capable of only binding the desired ion electric potential is generated by a separation of charge Membrane contains a ligand (L) that specifically and tightly binds analyte of interest (C+) The counter-ions (R-,A-) can’t cross the membrane and/or have low solubility in membrane or analyte solution A difference in the concentration of C+ exists across the outer membrane. 920316 slides

Ion-Selective Electrodes Potential across outer membrane depends on [C+] in analyte solution C+ diffuses across the membrane due to concentration gradient resulting in charge difference across membrane 920316 slides

Ion-Selective Electrodes A difference in the concentration of C+ exists across the inner membrane. C+ diffuses across the membrane due to concentration gradient resulting in charge difference across membrane Potential across inner membrane depends on [C+] in filling solution, which is a known constant Electrode potential is determined by the potential difference between the inner and outer membranes: 920316 slides

Ion-Selective Electrodes where Einner is a constant and Eouter depends on the concentration of C+ in analyte solution where [C+] is actually the activity of the analyte and n is the charge of the analyte 920316 slides

pH meter // ISE = (glass membrane ) that preferentially binds H+ 1) Combined (glass) electrod === ISE+ Ag/AgCl electrode 2) SCE outside electrod 2 Electrodes 920316 slides

pH Electrode pH Electrodes 1.) pH Measurement with a Glass Electrode Ag(s)|AgCl(s)|Cl-(aq) || H+(aq,outside) H+(aq,inside) , Cl-(aq)|AgCl(s)|Ag(s) Outer reference electrode [H+] outside (analyte solution) [H+] inside Inner reference electrode Eref1 Ej E1 E2 Eref2 Boundary potential difference (Eb) = E1 - E2 Glass membrane Selectively binds H+ Eb = c log[H+] Eb = c – pH 920316 slides Electric potential is generated by [H+] difference across glass membrane

iii. pH is determined by formation of boundary potential across glass membrane Boundary potential difference (Eb) = E1 - E2 where from Nernst Equation: Eb = c – 0.059pH Eb = c log[H+] -log aH+ (on exterior of probe or in analyte solution) constant Selective binding of cation (H+) to glass membrane 920316 slides

pH Electrodes Glass Membrane Irregular structure of silicate lattice
Cations (Na+) bind oxygen in SiO4 structure 920316 slides

pH Electrode pH Electrodes Glass Membrane Two surfaces of glass “swell” as they absorb water Surfaces are in contact with [H+] 920316 slides

pH Electrodes Glass Membrane
H+ diffuse into glass membrane and replace Na+ in hydrated gel region Ion-exchange equilibrium Selective for H+ because H+ is only ion that binds significantly to the hydrated gel layer Charge is slowly carried by migration of Na+ across glass membrane Potential is determined by external [H+] 920316 slides Constant and b are measured when electrode is calibrated with solution of known pH

iii. pH is determined by formation of boundary potential across glass membrane At each membrane-solvent interface, a small local potential develops due to the preferential adsorption of H+ onto the glass surface. Glass Surface 920316 slides

pH Electrode Junction Potential 1.) Occurs Whenever Dissimilar Electrolyte Solutions are in Contact Develops at solution interface (salt bridge) Small potential (few millivolts) Junction potential puts a fundamental limitation on the accuracy of direct potentiometric measurements Don’t know contribution to the measured voltage Different ion mobility results in separation in charge Again, an electric potential is generated by a separation of charge 920316 slides

Alkali Error H+ not only cation that can bind to glass surface - H+ generally has the strongest binding Get weak binding of Na+, K+, etc Most significant when [H+] or aH+ is low (high pH) - usually pH > 11-12 At low aH+ (high pH), amount of Na+ or K+ binding is significant  increases the “apparent” amount of bound H+ 920316 slides

Acid Error Errors at low pH (Acid error) can give readings that are too high Exact cause not known - usually occurs at pH > 0.5 Glass Electrodes for Other Cations change composition of glass membrane putting Al2O3 or B2O3 in glass enhances binding for ions other than H+ Used to make ISE’s for Na+, Li+, NH4+ 920316 slides

Redox titration curve 100 ml Fe M WITH Mno M (1M H2SO4) MnO4-+5Fe2++8H+  Mn2++5Fe3++4H20 Fe3+ + e Fe E0=0.771 5ml Before Eq 920316 slides

Attention 1!!!!!! 100 ml Fe M WITH Mno M MnO4-+5Fe2++8H+  Mn2++5Fe3++4H20 Fe2+  Fe3+ + e E0=-0.771 5ml Before Eq 920316 slides

Attention 2!!!!!! 100 ml Fe M WITH Mno M MnO4-+5Fe2++8H+  Mn2++5Fe3++4H20 5Fe2+  5Fe e E0=-0.771 5ml Before Eq 920316 slides

Titration curve 100 ml Fe M WITH Mno M MnO4-+5Fe2++8H+  Mn2++5Fe3++4H20 Fe3+ + e Fe E0=0.771 10 ml Before Eq 920316 slides

Titration curve 100 ml Fe M WITH Mno M MnO4-+5Fe2++8H+  Mn2++5Fe3++4H20 Fe3+ + e Fe E0=0.771 19.5 ml Before Eq 920316 slides

Titration curve 1e+Fe3+  Fe e+MnO4-+8H+  Mn2+ 20 ml At Eq ×5 920316 slides

Titration curve 5X Y 5Y X MnO4-+5Fe2++8H+  Mn2++5Fe3++4H20 /6 920316 slides

Titration curve 100 ml Fe M WITH Mno M MnO4-+5Fe2++8H+  Mn2++5Fe3++4H20 Fe3+  Fe2+ Mno4-+8H+  Mn2+ 20 ml At Eq 920316 slides

Keq MnO4-+5Fe2++8H+  Mn2++5Fe3++4H20 1e+Fe3+  Fe e+MnO4-+8H+  Mn2+ 20 ml At Eq 920316 slides

Keq At Eq 920316 slides

Keq MnO4-+5Fe2++8H+  Mn2++5Fe3++4H20 920316 slides

Keq MnO4-+5Fe2++8H+  Mn2++5Fe3++4H20 MnO4-+5e-+8H+→ Mn2++ 4H E0= n=5 Fe3+ +e → Fe E0= n=1 920316 slides

Titration curve 100 ml Fe M WITH Mno M MnO4-+5Fe2++8H+  Mn2++5Fe3++4H20 Mno4-+8H+  Mn2+ 20.5 ml After Eq 920316 slides

Titration curve 100 ml Fe M WITH Mno M MnO4-+5Fe2++8H+  Mn2++5Fe3++4H20 Mno4-+8H+  Mn2+ 21 ml After Eq 920316 slides

Titration curve ml of MnO4K E(v) ΔE/ΔV Δ2E/ΔV2 5 0.743 10 0.771 0.0056 15 0.799 19 0.846 19.5 0.866 0.04 0.0565 20 1.387 1.042 2.004 20.5 1.491 0.208 -1.668 21 1.494 0.006 -0.404 22 1.498 0.004 -0.002 23 1.5 0.002 25 1.503 0.0015 30 1.504 0.0002 920316 slides

Titration curve data Height is related to Keq Not related to concentration 920316 slides

Titration curve 920316 slides

Redox titrations & potentiometry

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