20 Applications of Oxidation/Reduction Titrations

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1 20 Applications of Oxidation/Reduction Titrations 20A AUXILIARY OXIDIZING AND REDUCING REAGENTS 20A1 Auxiliary Reducing Reagents Table 201 Uses of the Walden Reductor and the Jones Reductor Jones Walden Ag(s) + Cl AgCl(s) + e Zn(Hg)(s) Zn 2+ + Hg + 2e Fe 3+ + e Fe 2+ Fe 3+ + e Fe 2+ Cu 2+ + e Cu + Cu e Cu(s) H 2 MoO 4 + 2H + + e + MoO2 + 2H 2 O H 2 MoO 4 + 6H + + 3e Mo H 2 O UO H + + 2e U H 2 O UO H + + 2e U H 2 O UO H + + 3e U H 2 O V() H + + e VO H 2 O V() H + + 3e V H 2 O TiO 2+ not reduced TiO H + + e Ti H 2 O Cr 3+ not reduced Cr 3+ + e Cr 2+ 20A2 Auxiliary Oxidizing Reagents Sodium Bismuthate: Mn(II) MnO 4 NaBiO 3 (s) + 4H + + 2e BiO + + Na + + 2H 2 O Ammonium Peroxydisulfate, ammonium persulfate, (NH 4 ) 2 S 2 O 8 in acidic soln: Cr(III) dichromate Ce(III) Ce(IV) Mn(II) permanganate S 2 O e 2 2SO 4 Eº = 2.01 V The oxidations are catalyzed by traces of silver ion. The excess reagent is readily decomposed by a brief period of boiling: S 2 O H 2 O 4SO O 2 + 4H + Sodium Peroxide and Hydrogen Peroxide H 2 O 2 + 2H + + 2e 2H 2 O Eº = 1.78 V boiling: H 2 O 2 2H 2 O + 2O 2 (g) 20B APPLING STANDARD REDUCING AGENTS 20B1 Iron (II) Solutions iron(ii) ammonium sulfate, Fe(NH 4 ) 2 (SO 4 ) 2.6H 2 O (Mohr's salt) iron(ii) ethylenediamine sulfate, FeC 2 H 4 (NH 3 ) 2 (SO 4 ) 2.4H 2 O (Oesper's salt) Fig. 201 A Jones reductor. 137

2 Airoxidation of iron (II) takes place rapidly in neutral solutions but is inhibited in the presence of acids, with the most stable preparations being about 0.5 M in H 2 SO 4. oxidizing agents excess of standard Fe(II) standard soln of pot. dichromate or Ce(IV). Application: organic peroxides, hydroxylamine, Cr(VI), Ce(IV), Mo(VI), nitrate, chlorate, perchlorate and numerous other oxidants. 20B2 Sodium Thiosulfate iodine thiosulfate 2 2 2S 2 O 3 S 4 O 6 + 2e excess KI analyte/slightly acidic solution iodine standard solution of Na 2 S 2 O 3 ex: determination of sod. Hypochlorite in bleaches Ocl + 2I + 2H + Cl + I 2 + H 2 O (unmeasured excess KI) 2 I 2 + 2S 2 O 3 2I 2 + S 4 O 6 Detecting End Points in Iodine/Thiosulfate Titrations 1. disappearance of the iodine color M I2 discernible color 2. starch indicator deep blue color Starch undergoes decomposition in solution with high I 2 concentration. In titration of excess I 2 with Na 2 S 2 O 3, addition of the indicator must be deferred until most of the (b) I 2 has been reduced. Fig. 202 Thousands of glucose molecules polymerize to form huge molecules of β amylose as shown in (a). Molecules of βamylose tend to assume a helical structure. The iodine species I 3 as shown in (b) is incorporated into the amylose helix. The Stability of Sodium Thiosulfate Solutions decompose: S 2 O H + HSO 3 + S(s) ph, microorganisms, concentration of the solution, presence of Cu(II) ion and exposure to sunlight. Standardizing Thiosulfate Solutions primary standard: pot. iodate/excess KI (pot. dichromate, pot. bromate, pot. hydrogen iodate, pot. ferricyanide and metallic copper)/ excess KI. IO 3 + 5I + 6H + 3I 2 + 2H 2 O 1 mol IO 2 3 = 3 mol I 2 = 6 mol S 2 O 3 thiosulfate 138

3 Ex. 201 A solution of Sod. thiosulfate was standardized by dissolving g KIO 3 ( g/mol) in water, adding a large excess of KI, and acidifying with HCl. The liberated I 2 required ml of the thiosulfate soln to decolorize the blue starch/iodine complex. Calculate the molarity of the Na 2 S 2 O 3. amount Na C Na 2S2O3 2 S 2 O 3 = g KIO mmol = = M ml 3 1mol g or 6 = mmol 121mg mg/mmol = M ml Tab 202 Applications of Sodium Thiosulfate as Reductant Analyte HalfReaction Special Condition IO 4 IO 4 + 8H + + 7e ½ I 2 + 4H 2 O Acid solution IO 4 + 2H + + 2e IO 3 + H 2 O Neutral solution IO 3 IO 3 + 6H + + 5e ½ I 2 + 3H 2 O Strong acid BrO 3, ClO 3 XO 3 + 6H + + 6e X + 3H 2 O Strong acid Br 2, Cl 2 X 2 + 2I I 2 + 2X NO 2 HNO 2 + H + + e NO(g)+ H 2 O Cu 2+ Cu 2+ + I + e CuI(s) O 2 O 2 + 4Mn() 2 (s) + 2H 2 O 4Mn() 3 (s) Basic solution Mn() 3 (s) + 3H + + e Mn H 2 O Acidic solution O 3 O 3 (g) + 2H + + 2e O 2 (g) + H 2 O Organic peroxide RO + 2H + + 2e R + H 2 O 20C APPLYING STANDARD OXIDING AGENTS Table 203 Some common oxidants used as standard solutions Reagent and Reduction Standard Standardized Formula product Potential, V with Indicator* Stability# KMnO 4 Mn Na 2 C 2 O 4, Fe, As 2 O 3 MnO 4 (b) KBrO 3 Br 1.44 KBrO 3 (1) (a) Ce(IV), Ce 4+ Ce Na 2 C 2 O 4, Fe, As 2 O 3 (2) (a) K 2 Cr 2 O 7 Cr K 2 Cr 2 O 7, Fe (3) (a) I 2 I BaS 2 O 3 H 2 O, Na 2 S 2 O 3 starch (c) * (1) αnaphthoflavone; (2) 1,10phenanthroline iron(ii) complex (ferroin); (3) diphenylamine sulfonic acid. # (a) Indefinitely stable; (b) moderately stable, requires periodic standardization; (c) somewhat unstable, require frequent standardization. Eº' in H 2 SO

4 20C1 The Strong OxidantsPotassium Permanganate and Cerium(IV) MnO 4 + 8H + + 5e Mn H 2 O Eº = 1.51 V (in 0.1 M strong acid) Ce 4+ + e Ce 3+ Eº = 1.44 V (1M H 2 SO 4 ) = 1.70 V (1M HClO 4 ) = 1.61 V (1M HNO 3 ) Comparison of the Two Reagents Ce 4+ MnO 4 in sulfuric acid stable not oxidize Cl decompose slowly oxidize Cl HCl soln of analyte can be used cannot be used primarystandardgrade salt available selfindicator no color of MnO 4 cost (1L 0.02 M soln) $ 2.20 (4.40) $ 0.08 in < 0.1 M strong acid tendency to form ppt Detecting the End Points indicators: KMnO 4 solution intense purple color diphenylamine sulfonic acid 1, 10phenanthroline complex of Fe(II) 2MnO 4 + 3Mn 2+ + H 2 O 5MnO 2 (s) + 4H + K = equilibrium [MnO 4 ] rate: slow end point fades only gradually over 30s. in Ce(IV) titration: indicator: Fe(II) complex of 1,10phenanthroline or one of its substitute derivatives (Table 162) C 12 H 8 N 2 + Fe 2+ Fe(C 12 H 8 N 2 ) Fe(C 12 H 8 N 2 ) 3 + e Ferrous complex (Ferroin) red Ferric complex (Ferriin) weak blue The Preparation and Stability of Standard Solutions KMnO 4 soln: not entirely stable 4MnO 4 + 2H 2 O 4MnO 2 (s) + 3O 2 (g)+ 4 decomposition reaction is slow catalyzed by light, heat, acids, bases, Mn(II) and MnO

5 Ex. 202 Described how you would prepare 2.0 L of an approximately M soln of KMnO 4 ( g/mol). KMnO 4 needed = 2.0 L M g/mol = 3.16 g Dissolve about 3.2 g of KMnO 4 in a little water. After solution is complete add water to bring the volume to about 2.0 L. Heat the solution to boiling for a brief period, and let stand until it is cool. Filter through a glassfiltering crucible and stored in a clean dark bottle. Analytically Useful Cerium(IV) Compounds Name Formula Molar Mass Cerium(IV) ammonium nitrate Ce(NO 3 ) 4 2NH 4 NO Cerium(IV) ammonium sulfate Ce(SO 4 ) 2 2(NH 4 ) 2 SO 4 2H 2 O Cerium(IV) hydroxide Ce() Cerium(IV) hydrogen sulfate Ce(HSO 4 ) Primary Standards Sodium Oxalate. 2MnO 4 + 5H 2 C 2 O 4 + 6H + Mn(II) as a catalyst (autocatalysis) 2Ce 4+ + H 2 C 2 O 4 2Ce H + + 2CO 2 2Mn CO 2 (g) + 8H 2 O Ex. 203 You wish to standardize the soln in Ex.182 against pure Na 2 C 2 O 4 ( g/mol). If you want to use between 30 and 45 ml of the reagent for the standardization, what range of masses of the primary standard should you weigh out? for a 30mL titration: amount KMnO 4 = 30 ml M = 0.30 mmol mass Na 2 C 2 O 4 = 0.30 mmol 5/ = g for a 45mL titration: mass Na 2 C 2 O 4 = / = g Ex. 204 A g sample of primarystandard Na 2 C 2 O 4 required exactly ml of the KMnO 4 solution in Ex.302 to reach the end point. What was the molarity of the KMnO 4 reagent? amount Na 2 C 2 O 4 = g 1 mmol/0.134 g = mmol CKMnO 4 = mmol (2/5) (1/33.31 = M 141

6 Using Potassium Permanganate and Cerium(IV) Solutions: Table 205 Ex. 205 Aqueous solution containing approximately 3% (w/w) H 2 O 2 are sold in drug stores as a disinfectant. propose a method for determining the peroxide content of such a preparation using the standard soln described in Exs.183 and 4. Assume that you wish to use between 35 and 45 ml of the reagent for a titration. 5H 2 O 2 + 2MnO 4 + 6H + 5O 2 + 2Mn H 2 O ml reagent: amount KMnO 4 = (35 ~ 45) ml M = 0.401~0.515 mmol amount H 2 O 2 = (0.401 ~ 0.515) mmol (5/2) = 1.00~1.29 mmol mass sample = (1.00 ~ 1.29) (100/3) = 1.1 ~1.5 g Thus we could weigh out from 1.1 to 1.5 g samples. These should be diluted to perhaps 75 to 100 ml with water and made slightly acidic with dilute H 2 SO 4 before titration. 20C2 Potassium Dichromate Cr 2 O H + + 6e 2Cr H 2 O Eº = 1.33 V orange green in 1 M HCl or H 2 SO 4 Eº' = 1.0 ~ 1.1 V Advantages: stable, can be boiled without decomposition and do not react with HCl, primarystandard reagent is available and at a modest cost. Disadvantage: lower electrode potential and the slowness reaction. Preparing Dichromate Solutions reagentgrade K 2 Cr 2 O 7 dried at 150 to 200 before being weighed indicator: diphenylamine sulfonic acid, violet (oxidized) colorless (reduced) Appling Potassium Dichromate Solutions 1. titration of Fe(II): in moderate conc. of HCl Cr 2 O Fe H + 2Cr Fe H 2 O 2. indirect determination of oxidizing agents (nitrate, chlorate, permanganate, dichromate and organic peroxides): analyte/acidic solution + measured excess Fe(II) backtitrated excess Fe(II) 142

7 Ex. 206 A 5.00mL sample of brandy was diluted to L in a volumetric flask. The ethanol(c 2 H 5 ) in a 25.00mL aliquot of the diluted soln was distilled into ml of M K 2 Cr 2 O 7, and oxidized to acidic acid with heating. 3C 2 H 5 + 2Cr 2 O H + 4Cr CH 3 CO + 11H 2 O After cooling, ml of M Fe 2+ were pipetted into the flask. The excess Fe 2+ was then titrated with 7.46 ml of the standard K 2 Cr 2 O 7 to a diphenylamine sulfonic acid end point. Calculate the percent (w/v) C 2 H 5 (46.07 g/mol) in the brandy. amount K 2 Cr 2 O 7 = ( ) ml = mmol K 2 Cr 2 O 7 consumed by Fe 2+ = /6 = mmol K 2 Cr 2 O 7 consumed by C 2 H 5 = = mmol mass C 2 H 5 = (3/2) = g percent C 2 H 5 = /( /1000) 100 % = % 20C3 Iodine weak oxidizing agents: determination of strong reductants I 3 + 2e 3I Eº = V advantages: selectivity, sensitive and reversible indicator disadvantage: lack stability Properties of Iodine Solutions I 2 (s) + I I 3 K = lack stability: volatility of iodine, slowly attacks most organic materials, airoxidation of iodide ion ( conc.). 4I + O 2 (g) + 4H + 2I 2 + 2H 2 O Standardizing and Appling Iodine Solutions Standardization: anhydrous Na thiosulfate or Ba thiosulfate Iodimetry: (direct method) I 2 reducing agents (ex: thiosulfate or arsenites) Iodometry: (indirect method) oxidizing agents + excess KI I 2 thiosulfate Indicator: 1. Starch indicator solution I 2 + I I 3 + starch I 3 starch complex (bluepurple) 2. CCl 4, HCCl 3, CS 2 I 2 in CCl 4, HCCl 3, CS 2 violet color 143

8 Preparation of 0.1 N Iodine solution 12.7 g I g KI/20 ml H 2 O adding H 2 O to 1 L Standardization Primary standard: Arsenic (III) oxide, As 2 O 3 As 2 O 3 + 6Na 2Na 3 AsO 3 + 3H 2 O Na 3 AsO 3 + 3HCl H 3 AsO 3 + 3NaCl H 3 AsO 3 + H 2 O + I 2 H 3 AsO 4 + 2HI ph : H 3 AsO 3 + H 2 O + I 2 H 3 AsO 4 + 2H + + 2I ph : I IO + I + H 2 O 3IO IO 3 + 2I strong oxidizing agent 滴定中加入 NaHCO 3 [ph: 7~8] Na 3 AsO 3 + I 2 + 2NaHCO 3 Na 3 AsO 4 + 2NaI + 2CO 2 + H 2 O mg As2O3 Calculation: N of Iodine = / 4 ml Iodine (1) Direct Iodimetric Titration a. H 3 AsO 3 + H 2 O + I 2 H 3 AsO 4 + 2H + + 2I b. Assay of Ascorbic Acid (Vit C) O HO O (enediol) CH CH 2 O O CH CH 2 + I 2 O O (αdiketone) (2) Residual Titration (I 2 Na 2 S 2 O 3 ) NaHSO 3 + I 2 +H 2 O NaHSO 4 + 2HI I Na 2 S 2 O 3 2NaI + Na 2 S 4 O 6 (3) Iodometry: Sample + KI (excess) I2 Na2S2O3 + 2I + 2H+ (+2) (1e) (+1) a. CuSO 4 + 4KI 2CuI + I 2 + 2K 2 SO 4 I 2 + 2Na 2 S 2 O 3 2NaI + Na 2 S 4 O 6 144

9 b. Assay of sodium hypochlorite solution (NaOCl) NaOCl + H + HOCl (+1) (2e) (1) HOCl + 2KI + HOAc I 2 + KCl + KOAc + H 2 O I 2 + 2Na 2 S 2 O 3 2NaI + Na 2 S 4 O 6 Table 206 Some Applications of Iodine Solutions Analyte HalfReaction As H 3 ASO 3 + H 2 O H 3 AsO 4 + 2H + + 2e Sb H 3 SbO 3 + H 2 O H 3 SbO 4 + 2H + + 2e Sn Sn 2+ Sn e H 2 S H 2 S S(s) + 2H + + 2e SO 2 SO H 2 O SO H + + 2e 2 S 2 O S 2 O 3 S 4 O e N 2 H 4 N 2 H 4 N 2 (g) + 4H + + 4e Ascorbic acid C 6 H 8 O 6 C 6 H 6 O 6 + 2H + + 2e ** Dichloroindophenol Titration Determination of Ascorbic acid Preparation O N Cl HO NH Cl Cl Cl Blue in basic sol'n colorless Pink in acid sol'n End point: pink (selfindicator) Vit C titration in metaphosphoric acid and acetic acid sol'n 20C4 Potassium Bromate as a Source of Bromine Primarystandard KBrO 3 is available, stable standard 0.1 N Bromine sol'n (Koppeschaar's sol'n): (3 g KBrO g KBr)/1 L H 2 O Assay of sample: aniline, phenol, salicylic acid, resorcinol etc. BrO 3 + 5Br + 6H + 3Br 2 + 3H 2 O standard soln excess 1 mol KBrO 3 = 3 mol Br 2 2I + Br 2 I 2 + 2Br 2 I 2 + 2S 2 O 3 S 4 O I (excess KI) 145

10 Br Br + 3 Br H + + 3Br Br (ppt) soluble in CHCl 3 CO Br Br + 3Br 2 + HBr + CO 2 Br Substitution Reactions halogen substitution: replacement of H in an aromatic ring by a halogen. determination of aromatic compound that contain strong orthoparadirecting groups, particularly amines and phenols. Ex: 1. Determination of 8hydroxyquinoline N + 2Br 2 2. Determination of aluminum Al 3+ N ph 49 Br Br + 2HBr + ( OC H N) ( S ) HOC H N Al H 9 6 hot 4M HCl 3+ ( OC H N) ( ) 3HOC H N + Al S Al 9 H6N + 6Br2 3HOC9H4NBr + 6HBr 3HOC 2 Ex. 207 A g sample of an antibiotic powder containing sulfanilamide was dissolved in HCl and the solution diluted to ml. A 20.00mL aliquot was transferred to a stoppered flask and ml of M KBrO 3 added. About 10 g of KBr was added to form Br 2, which brominated the sulfanilamide in the sample. After 10 min, an excess of KI was added and the liberated iodine titrated with ml of M sodium thiosulfate. The reaction are BrO 3 + 5Br + 6H + 3Br 2 + 3H 2 O NH 2 NH 2 Br Br + 2Br 2 + 2H + + 2Br SO 2 NH 2 SO 2 NH 2 146

11 Br 2 + 2I 2Br + I 2 (excess KI) I 2 + 2S 2 O 3 2 S 4 O I Calculate the % NH 2 C 6 H 4 SO 2 NH 2 ( g/mol) in the powder. total amount Br 2 = ml M 3 = mmol amount excess Br 2 = amount I 2 = ml M (1/2) = mmol The amount of Br 2 consumed by the sample = = mmol mass analyte = (1/2) = g % analyte = 100% = 80.47% 20.00mL mL ( ) 5 or 4 100% = 80.47% Addition Reactions H C C H H H + Br 2 H C C Br Br 20C5 Determining Water with the Karl Fischer Reagent H Karl Fischer Reagent: I 2, SO 2, organic base such as pyridine (C 5 H 5 N) or imidazole /CH 3 or lowmolecularmass alcohol In aprotic solvent: I 2 + SO 2 + 2H 2 O 2HI + H 2 SO 4 2 mol H 2 O 1 mol I 2 Classical chemistry: use anhydrous methanol as solvent, and excess pyridine C 5 H 5 N I 2 + C 5 H 5 N SO 2 + C 5 H 5 N + H 2 O 2C 5 H 5 N HI + C 5 H 5 N SO 3 C 5 H 5 N + SO 3 + CH 3 C 5 H 5 N(H)SO 4 CH 3 1 mol H 2 O 1 mol I 2, 1 mol SO 2, 3 mol C 5 H 5 N Pyridinefree chemistry Replaced by other amines: imidazole (1) Solvolysis: 2R + SO 2 RSO R 2 (2) Buffering: B + RSO 3 + R + 2 BH + SO 3 R + R (3) Redox: B I 2 + BH + SO 3 R + B + H 2 O BH + SO 4 R + 2BH + I 1 mol H 2 O 1 mol I 2 Interfering reactions 147