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920316http:\asadipour.kmu.ac.ir 81 slides 1 Redox titrations & potentiometry (Mark=3)

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2 920316http:\asadipour.kmu.ac.ir 81 slides 1 Redox titrations & potentiometry (Mark=3)

3 Redox reactions half-reactions: Reduction 2Fe e -  2Fe 2+ oxidation Sn 2+  Sn e - 2Fe 3+ + Sn 2+  2Fe 2+ + Sn http:\asadipour.kmu.ac.ir 81 slides

4 Standard Reduction Potentials (strength) Reduction Half-Reaction E  (V) F 2 (g) + 2e -  2F - (aq) 2.87 Au 3+ (aq) + 3e -  Au(s) 1.50 Cl 2 (g) + 2 e -  2Cl - (aq) 1.36 Cr 2 O 7 2- (aq) + 14H + (aq) + 6e -  2Cr 3+ (aq) + 7H 2 O 1.33 O 2 (g) + 4H + + 4e -  2H 2 O(l) 1.23 Ag + (aq) + e -  Ag(s) 0.80 Fe 3+ (aq) + e -  Fe 2+ (aq) 0.77 Cu 2+ (aq) + 2e -  Cu(s) 0.34 Sn 4+ (aq) + 2e -  Sn 2+ (aq) H + (aq) + 2e -  H 2 (g) 0.00 Sn 2+ (aq) + 2e -  Sn(s) Ni 2+ (aq) + 2e -  Ni(s) Fe 2+ (aq) + 2e -  Fe(s) Zn 2+ (aq) + 2e -  Zn(s) Al 3+ (aq) + 3e -  Al(s) Mg 2+ (aq) + 2e -  Mg(s) Li + (aq) + e -  Li(s) Ox. agent strength increases Red. agent strength increases http:\asadipour.kmu.ac.ir 81 slides

5 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 H 2 O 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 http:\asadipour.kmu.ac.ir 81 slides

6 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 H 2 O 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 http:\asadipour.kmu.ac.ir 81 slides

7 Balancing of redox reactions H 2 O 2 (aq) + Cr 2 O 7 -2 (aq )  Cr 3+ (aq) + O 2 (g) Redox reaction ====================================== 1)write 2 half reactions Half Rxn (red): Cr 2 O 7 -2 (aq)  Cr 3+ Half Rxn (oxid): H 2 O 2 (aq)  O 2 2)Atom balance Cr 2 O 7 -2 (aq)  2Cr 3+ H 2 O 2 (aq)  O http:\asadipour.kmu.ac.ir 81 slides

8 Balancing of redox reactions 3)Oxygen balance Half Rxn (red): Cr 2 O 7 -2 (aq)  2Cr H 2 O Half Rxn (oxi): H 2 O 2 (aq)  O 2 4)Hydrogen balance Half Rxn (red): 14H + + Cr 2 O 7 -2 (aq)  2Cr H 2 O Half Rxn (oxi): H 2 O 2 (aq)  O 2 + 2H + 5)Electron balance 6e H + + Cr 2 O 7 -2 (aq)  2Cr H 2 O H 2 O 2 (aq)  O 2 + 2H + + 2e http:\asadipour.kmu.ac.ir 81 slides

9 Balancing of redox reactions 6) Equalize of produced and consumed electrons 6e H + + Cr 2 O 7 -2 (aq)  2Cr H 2 O ( H 2 O 2 (aq)  O 2 + 2H + + 2e - ) x 3 7)Multiply each half reaction 8 H + + 3H 2 O 2 + Cr 2 O 7 2-  2Cr O 2 + 7H 2 O http:\asadipour.kmu.ac.ir 81 slides

10 Balance the redox reactions I 2 +S 2 O 3 2- ⇋ I - +S 4 O 6 2- I 2 +S 2 O 3 2- ⇋ I - +SO http:\asadipour.kmu.ac.ir 81 slides H+H+ OH -

11 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 Ce 4+ Preoxidation : Peroxydisulfate ( (NH 4 ) 2 S 2 O 8 ) 2– ) Sodium bismuthate ( NaBiO 3 ) Hydrogen peroxide (H2O2) slides Prereduction : Stannous chloride ( SnCl 2 ) Chromous chloride Jones reductor (zinc coated with zinc amalgam) Walden reductor ( solid Ag and 1M HCl)

12 Redox titrations 1)In solution (visual indicators) 2Fe 3+ + Sn 2+  2Fe 2+ + Sn http:\asadipour.kmu.ac.ir 81 slides 2) In electrochemical cell. (Potentiometry)

13 http:\asadipour.kmu.ac.ir 81 slides Oxidizing and reducing agents

14 http:\asadipour.kmu.ac.ir 81 slides

15 Iodimetry and iodometry Iodimetry: 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 http:\asadipour.kmu.ac.ir 81 slides I - + Cu 2+ → I 2 + Cu + I 2 + S 2 O 3 2- → 2I - + S 4 O 6 2-

16 1) Iodine only dissolves slightly in water. Its solubility is enhanced by interacting with I - Preparation of aqoues solution I http:\asadipour.kmu.ac.ir 81 slides Standardization of Iodin with Arsenious oxide, As 2 O 3 + 3H 2 O = 2H 3 AsO 3 As 4 O 6 + 6H 2 O = 4H 3 AsO 3 H 3 AsO 3 + I 3 – + H 2 O = H 3 AsO 4 + 3I – + 2H +

17 1) An excellent way to prepare standard I 3 - 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 I 3 - by quantitative reverse disproportionation: 2) Cu + HNO 3  Cu 2+ Cu 2+ +4I -  2CUI + I 2 standard I http:\asadipour.kmu.ac.ir 81 slides

18 Stability of I 2 Solutions In acidic solutions of I 3 - 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 http:\asadipour.kmu.ac.ir 81 slides I 2 + OH - ⇌ IO - + I - + H + 3IO - ⇌ IO I -

19 http:\asadipour.kmu.ac.ir 81 slides Iodimetry

20 http:\asadipour.kmu.ac.ir 81 slides iodometry

21 Sodium thiosulfate, Na 2 S 2 O 3 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: 2S 2 O 3 2–  S 4 O 6 2– + 2e E o = 0.08 Ex. Determination of hypochlorite in bleaches [CaCl(OCl)H 2 O]: OCl – + 2I – + 2H +  Cl – + I 2 + H 2 O (unmeasured excess KI) I S 2 O 3 2–  2I – + S 4 O 6 2– Indicator: soluble starch (  -amylose) http:\asadipour.kmu.ac.ir 81 slides

22 Standardization of thiosulfate solution: Primary standard : potassium iodate (KIO 3 ), K 2 Cr 2 O 7, KBrO 3 Titration reactions: KIO 3 + 5KI + 6HCl  3I 2 + 6KCl + 3 H 2 O I 2 + 2Na 2 S 2 O 3  2NaI + Na 2 S 4 O 6 KIO 3  3I 2  6Na 2 S 2 O 3 ·5H 2 O  6 Equivalent S 2 O H + ⇋ HSO 3 - +S (s) pH, Microorganisms, Concentration, Cu 2+, Sunlight Stabilizer for sodium thiosulfate solution : Na 2 CO 3 Na 2 S 2 O 3 + H 2 O + CO 2  Na 2 CO 3 + H 2 S 2 O 3 H 2 S 2 O 3  H 2 SO 3 + S http:\asadipour.kmu.ac.ir 81 slides

23 slides

24 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 V to V with respect SHE http:\asadipour.kmu.ac.ir 81 slides

25 http:\asadipour.kmu.ac.ir 81 slides

26 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 I 2, not to a change in redox potential. Starch-Iodine Complex Structure of the repeating unit of the sugar amylose http:\asadipour.kmu.ac.ir 81 slides

27 Permanganate titration Oxidation with permanganate : (Reduction of permanaganate) KMnO 4 Powerful oxidant that the most widely used. 1) In strongly acidic solutions (1M H 2 SO 4 or HCl, pH  1) MnO 4 – + 8H + + 5e = Mn H 2 O E o = 1.51 V KMnO 4 is a self-indicator. 2) In feebly acidic, neutral, or alkaline solutions MnO 4 – + 4H + + 3e = MnO 2 (s) + 2H 2 O E o = V 3) In very strongly alkaline solution (2M NaOH) MnO 4 – + e = MnO 4 2 – E o = V  http:\asadipour.kmu.ac.ir 81 slides

28 Permanganate titration Duration of colour in end point (30 seconds) MnO 4 – + 3Mn H 2 O  5MnO 2 + 4H + K=1*10 47 Stability of aqoues solution of MnO 4 - MnO 4 – + 2H 2 O  4MnO 2 (s) + 3O 2 (g) +4OH -  http:\asadipour.kmu.ac.ir 81 slides

29 Standardization of KMnO 4 solution Potassium permanganate is not primary standard, because traces of MnO 2 are invariably present. Standardization by titration of sodium oxalate (primary standard) : 2KMnO Na 2 (COO) 2 + 8H 2 SO 4 = 2MnSO 4 + K 2 SO 4 + 5Na 2 SO CO 2 + 8H 2 O 2KMnO 4  5 Na 2 (COO) 2  10 Equivalent http:\asadipour.kmu.ac.ir 81 slides

30 Preparation of 0.1 N potassium permanganate solution KMnO 4 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 KMnO 4 (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. 2)Filter the liquid through a sintered-glass filter crucible to remove solid MnO 2. 3)Transfer the filtrate to a clean stoppered bottle freed from grease with cleaning mixture and standardize it. 4) Protect the solution from evaporation, dust, and reducing vapors, and keep it in the dark or in diffuse light. 4)If in time managanese dioxide settles out, refilter the solution and restandardize it http:\asadipour.kmu.ac.ir 81 slides

31 Applications of permanganometry (1)H 2 O 2 2KMnO H 2 O 2 + 3H 2 SO 4 = 2MnSO 4 + K 2 SO 4 + 5O 2 + 8H 2 O (2) NaNO 2 2NaNO 2 + H 2 SO 4 = Na 2 SO 4 + HNO 2 2KMnO HNO 2 + 3H 2 SO 4 = 2MnSO 4 + K 2 SO 4 + 5HNO 3 + 3H 2 O (3) FeSO 4 2KMnO FeSO 4 + 8H 2 SO 4 = 2MnSO 4 + K 2 SO 4 + 5Fe 2 (SO 4 ) 3 + 8H 2 O (4) CaO CaO + 2HCl = CaCl 2 + H 2 O CaCl 2 + H 2 C 2 O 4 = CaC 2 O 4 + 2HCl (excess oxalic acid) 2KMnO H 2 C 2 O 4 + 3H 2 SO 4 = 2MnSO 4 + K 2 SO CO 2 + 8H 2 O (back tit) (5) Calcium gluconate [CH 2 OH(CHOH) 4 COO] 2 Ca + 2HCl = CaCl 2 + 2CH 2 OH9CHOH) 4 COOH (NH 4 ) 2 C 2 O 4 + CaCl 2 = CaC 2 O NH 4 Cl CaCl 2 + H 2 SO 4 = H 2 C 2 O 4 + CaSO 4 2KMnO H 2 C 2 O 4 + 3H 2 SO 4 = 2MnSO 4 + K 2 SO CO 2 + 8H 2 O http:\asadipour.kmu.ac.ir 81 slides

32 Bromatimetry BrO 3 – + 5Br – + 6H +  3Br 2 + H 2 O 2I – + Br 2  I 2 + 2Br – http:\asadipour.kmu.ac.ir 81 slides I S 2 O 3 2–  2I – + S 4 O 6 2–

33 Determining water with the Karl Fisher Reagent The Karl Fisher reaction : I 2 + SO 2 + 2H 2 O  2HI + H 2 SO 4 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 : C 5 H 5 NI 2 + C 5 H 5 NSO 2 + C 5 H 5 N + H 2 O  2 C 5 H 5 NHI + C 5 H 5 NSO 3 C 5 H 5 N + SO 3 – + CH 3 OH  C 5 H 5 N(H)SO 4 CH 3 Pyridinium sulfite can also consume water. C 5 H 5 N + SO 3 – + H 2 O  C 5 H 5 NH + SO 4 H – 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 http:\asadipour.kmu.ac.ir 81 slides

34 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 http:\asadipour.kmu.ac.ir 81 slides

35 Electrodes and Potentiometry  Potential change only dependent on one ½ cell concentrations  Reference electrode is fixed or saturated  doesn’t change! Reference electrode, [Cl - ] is constant Potential of the cell only depends on [Fe 2+ ] & [Fe 3+ ] Pt wire is indicator electrode whose potential responds to [Fe 2+ ]/[Fe 3+ ] Unknown solution of [Fe 2+ ] & [Fe 3+ ] slides E cell =E cathod -E anod Anod is conventionally reference electode Fe 3+ +e -  Fe 2+ AgCl (s) + e - → Ag + Cl -

36 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 H 2 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 http:\asadipour.kmu.ac.ir 81 slides > > >

37 Reference Electrodes 1.)Silver-Silver Chloride Reference Electrode E o = V Activity of Cl - not 1  E (sat,KCl) = V http:\asadipour.kmu.ac.ir 81 slides

38 Reference Electrodes 2.)Saturated Calomel Reference Electrode (S.C.E)  Saturated KCl maintains constant [Cl - ] even with some evaporation E o = V Activity of Cl - not 1  E (sat,KCl) = V http:\asadipour.kmu.ac.ir 81 slides

39 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 http:\asadipour.kmu.ac.ir 81 slides

40 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 a) Metallic Electrodes of the First Kind i. Involves single reaction ii. Detection of cathione derived from the metal used in the electrode iii. Example: use of copper electrode to detect Cu 2+ in solution ½ reaction: Cu e -  Cu (s) E ind gives direct measure of Cu 2+ : E ind = E o Cu – (0.0592/2) log a Cu(s) /a Cu 2+ since a Cu(s) = 1: E ind = E o Cu – (0.0592/2) log 1/a Cu 2+ or using pCu = -log a Cu 2+ : E ind = E o Cu – (0.0592/2) pCu http:\asadipour.kmu.ac.ir 81 slides

41 b) Metallic Electrodes of the Second Kind i. Detection of anion derived from the interaction with metal ion (M n+ ) from the electrode ii. Anion forms precipitate or stable complex with metal ion (M n+ ) iii. Example: Detection of Cl - with Ag electrode ½ reaction: AgCl(s) + e -  Ag(s) + Cl - E O = V E ind gives direct measure of Cl - : E ind = E o – (0.0592/1) log a Ag(s) a Cl - /a AgCl(s) since a Ag(s) and a AgCl(s) = 1 & E o = V: E ind = – (0.0592/1) log a Cl http:\asadipour.kmu.ac.ir 81 slides

42 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 Ce 3+ with Pt electrode ½ reaction: Ce 4+ + e -  Ce 3+ E ind responds to Ce 4+ : E ind = E o – (0.0592/1) log a Ce 3+ /a Ce http:\asadipour.kmu.ac.ir 81 slides

43 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 http:\asadipour.kmu.ac.ir 81 slides

44 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 E=E1-E2 Constant C + - membrane C + - Un known solution Constant C + - Inner solution E2 E1 81 slides ISE + IRE = Combined electrod

45 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 Membrane contains a ligand (L) that specifically and tightly binds analyte of interest (C + ) A difference in the concentration of C + exists across the outer membrane. The counter-ions (R -,A - ) can’t cross the membrane and/or have low solubility in membrane or analyte solution http:\asadipour.kmu.ac.ir 81 slides electric potential is generated by a separation of charge

46 C + diffuses across the membrane due to concentration gradient resulting in charge difference across membrane Ion-Selective Electrodes Potential across outer membrane depends on [C+] in analyte solution http:\asadipour.kmu.ac.ir 81 slides

47 Ion-Selective Electrodes Potential across inner membrane depends on [C+] in filling solution, which is a known constant C + diffuses across the membrane due to concentration gradient resulting in charge difference across membrane A difference in the concentration of C + exists across the inner membrane. Electrode potential is determined by the potential difference between the inner and outer membranes: http:\asadipour.kmu.ac.ir 81 slides

48 Ion-Selective Electrodes where E inner is a constant and E outer depends on the concentration of C + in analyte solution http:\asadipour.kmu.ac.ir 81 slides where [C + ] is actually the activity of the analyte and n is the charge of the analyte

49 pH meter // ISE = (glass membrane ) that preferentially binds H + 1) Combined (glass) electrod === ISE + Ag/AgCl electrode 2) SCE outside electrod http:\asadipour.kmu.ac.ir 81 slides 2 Electrodes

50 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 Glass membrane Selectively binds H + Electric potential is generated by [H + ] difference across glass membrane http:\asadipour.kmu.ac.ir 81 slides E ref1 E1E1 E2E2 EjEj E ref2 Boundary potential difference (E b ) = E 1 - E 2 E b = c log[H + ] E b = c – pH

51 iii. pH is determined by formation of boundary potential across glass membrane Boundary potential difference (E b ) = E 1 - E 2 where from Nernst Equation: E b = c – 0.059pH -log a H + (on exterior of probe or in analyte solution) constant Selective binding of cation (H + ) to glass membrane http:\asadipour.kmu.ac.ir 81 slides E b = c log[H + ]

52 pH Electrode pH Electrodes Glass Membrane  Irregular structure of silicate lattice Cations (Na + ) bind oxygen in SiO 4 structure http:\asadipour.kmu.ac.ir 81 slides

53 pH Electrode pH Electrodes Glass Membrane  Two surfaces of glass “swell” as they absorb water - Surfaces are in contact with [H + ] http:\asadipour.kmu.ac.ir 81 slides

54 pH Electrode 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 + ] Constant and b are measured when electrode is calibrated with solution of known pH http:\asadipour.kmu.ac.ir 81 slides

55 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 http:\asadipour.kmu.ac.ir 81 slides

56 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 Again, an electric potential is generated by a separation of charge Different ion mobility results in separation in charge http:\asadipour.kmu.ac.ir 81 slides

57 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 a H + is low (high pH) - usually pH > At low a H + (high pH), amount of Na + or K + binding is significant  increases the “apparent” amount of bound H http:\asadipour.kmu.ac.ir 81 slides

58 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 Al 2 O 3 or B 2 O 3 in glass enhances binding for ions other than H + Used to make ISE’s for Na +, Li +, NH http:\asadipour.kmu.ac.ir 81 slides

59 Redox titration curve 100 ml Fe M WITH Mno M (1M H 2 SO 4 ) MnO Fe 2+ +8H +  Mn 2+ +5Fe 3+ +4H 2 0 Fe 3+ + e  Fe 2+ E 0 = ml Before Eq http:\asadipour.kmu.ac.ir 81 slides

60 Attention 1!!!!!! 100 ml Fe M WITH Mno M MnO Fe 2+ +8H +  Mn 2+ +5Fe 3+ +4H 2 0 Fe 2+  Fe 3+ + e E 0 = ml Before Eq http:\asadipour.kmu.ac.ir 81 slides

61 Attention 2!!!!!! 100 ml Fe M WITH Mno M MnO Fe 2+ +8H +  Mn 2+ +5Fe 3+ +4H 2 0 5Fe 2+  5Fe e E 0 = ml Before Eq http:\asadipour.kmu.ac.ir 81 slides

62 Titration curve 100 ml Fe M WITH Mno M MnO Fe 2+ +8H +  Mn 2+ +5Fe 3+ +4H 2 0 Fe 3+ + e  Fe 2+ E 0 = ml Before Eq http:\asadipour.kmu.ac.ir 81 slides

63 Titration curve 100 ml Fe M WITH Mno M MnO Fe 2+ +8H +  Mn 2+ +5Fe 3+ +4H 2 0 Fe 3+ + e  Fe 2+ E 0 = ml Before Eq http:\asadipour.kmu.ac.ir 81 slides

64 Titration curve 100 ml Fe M WITH Mno M MnO Fe 2+ +8H +  Mn 2+ +5Fe 3+ +4H 2 0 Fe 3+ + e  Fe 2+ E 0 = ml Before Eq http:\asadipour.kmu.ac.ir 81 slides

65 Titration curve 100 ml Fe M WITH Mno M MnO Fe 2+ +8H +  Mn 2+ +5Fe 3+ +4H 2 0 Fe 3+ + e  Fe 2+ E 0 = ml Before Eq http:\asadipour.kmu.ac.ir 81 slides

66 Titration curve 1e+Fe 3+  Fe 2+ 5e+MnO H +  Mn ml At Eq http:\asadipour.kmu.ac.ir 81 slides ×5

67 Titration curve http:\asadipour.kmu.ac.ir 81 slides MnO Fe 2+ +8H +  Mn 2+ +5Fe 3+ +4H 2 0 /6 5X X Y 5Y

68 Titration curve 100 ml Fe M WITH Mno M MnO Fe 2+ +8H +  Mn 2+ +5Fe 3+ +4H 2 0 Fe 3+  Fe 2+ Mno H +  Mn ml At Eq http:\asadipour.kmu.ac.ir 81 slides

69 Keq 1e+Fe 3+  Fe 2+ 5e+MnO H +  Mn ml At Eq http:\asadipour.kmu.ac.ir 81 slides MnO Fe 2+ +8H +  Mn 2+ +5Fe 3+ +4H 2 0

70 Keq http:\asadipour.kmu.ac.ir 81 slides At Eq

71 Keq http:\asadipour.kmu.ac.ir 81 slides MnO Fe 2+ +8H +  Mn 2+ +5Fe 3+ +4H 2 0

72 K eq MnO Fe 2+ +8H +  Mn 2+ +5Fe 3+ +4H 2 0 MnO e - +8H + → Mn H 2 0 E 0 =1.51 n=5 Fe 3+ +e → Fe 2 + E 0 = n= http:\asadipour.kmu.ac.ir 81 slides

73 Titration curve 100 ml Fe M WITH Mno M MnO Fe 2+ +8H +  Mn 2+ +5Fe 3+ +4H 2 0 Mno H +  Mn ml After Eq http:\asadipour.kmu.ac.ir 81 slides

74 Titration curve 100 ml Fe M WITH Mno M MnO Fe 2+ +8H +  Mn 2+ +5Fe 3+ +4H 2 0 Mno H +  Mn ml After Eq http:\asadipour.kmu.ac.ir 81 slides

75 Titration curve 100 ml Fe M WITH Mno M MnO Fe 2+ +8H +  Mn 2+ +5Fe 3+ +4H 2 0 Mno H +  Mn ml After Eq slides

76 Titration curve 100 ml Fe M WITH Mno M MnO Fe 2+ +8H +  Mn 2+ +5Fe 3+ +4H 2 0 Mno H +  Mn ml After Eq slides

77 Titration curve 100 ml Fe M WITH Mno M MnO Fe 2+ +8H +  Mn 2+ +5Fe 3+ +4H 2 0 Mno H +  Mn ml After Eq http:\asadipour.kmu.ac.ir 81 slides

78 Titration curve 100 ml Fe M WITH Mno M MnO Fe 2+ +8H +  Mn 2+ +5Fe 3+ +4H 2 0 Mno H +  Mn ml After Eq http:\asadipour.kmu.ac.ir 81 slides

79 Titration curve http:\asadipour.kmu.ac.ir 81 slides ml of MnO4K E(v) Δ E/ Δ V Δ2 E/ Δ V

80 Titration curve data http:\asadipour.kmu.ac.ir 81 slides Height is related to K eq Not related to concentration

81 Titration curve http:\asadipour.kmu.ac.ir 81 slides

82 Titration curve http:\asadipour.kmu.ac.ir 81 slides


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