ndirect EDTA determination of cesium. Cesium ion does not form a strong EDTA complex, but it can be analyzed by adding a known volume of NaBil, in cold concentrated acetic acid containing excess Nal. Solid Cs, Bi, I, is precipitated, filtered, and emoved. The excess yellow Bil, is then titrated with EDTA. The end point occurs when the yellow color disappears. (Sodium hiosulfate is used in the reaction to prevent the liberated I from being oxidized to yellow aqueous I, by O, from the air.) The orecipitation is fairly selective for Cst. The ions Lit, Na*, K*, and low concentrations of Rbt do not interfere, although TIt loes. Suppose that 23.00 mL of an unknown containing Cs+ were treated with 23.00 mL of 0.07719 M NaBil, and the inreacted Bil, required 18.74 mL of 0.0552 M EDTA for a complete titration. Find the concentration of Cs+ in the unknown. [Cs*] =

ndirect EDTA determination of cesium. Cesium ion does not form a strong EDTA complex, but it can be analyzed by adding a known volume of NaBil, in cold concentrated acetic acid containing excess Nal. Solid Cs, Bi, I, is precipitated, filtered, and emoved. The excess yellow Bil, is then titrated with EDTA. The end point occurs when the yellow color disappears. (Sodium hiosulfate is used in the reaction to prevent the liberated I from being oxidized to yellow aqueous I, by O, from the air.) The orecipitation is fairly selective for Cst. The ions Lit, Na, K, and low concentrations of Rbt do not interfere, although TIt loes. Suppose that 23.00 mL of an unknown containing Cs+ were treated with 23.00 mL of 0.07719 M NaBil, and the inreacted Bil, required 18.74 mL of 0.0552 M EDTA for a complete titration. Find the concentration of Cs+ in the unknown. [Cs*] =

Fundamentals Of Analytical Chemistry

9th Edition

ISBN:9781285640686

Author:Skoog

Publisher:Skoog

Chapter31: Introduction To Analytical Separations

Section: Chapter Questions

Problem 31.18QAP

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a 10 ml sample containing fe3+ and cu2+ was diluted to 100.0 mL solution. A 20 mL aliquot required 25 mL of 0.00500 M EDTA for complete titration. A separate 20 mL aliquot of the diluted sample was treated with NH4F to protect the Fe3+. Then the Cu2+ was treated with thiourea. Upon addition of 25.0 ml of 0.00500 M EDTA, the Fe3+ was liberated from its fluoride complex and formed an EDTA complex. The excess EDTA required 21.5 ml of 0.0180 M Pb2+ to reach an endpoint, using xylenol orange. (a) identify the masking agent (b) identify the auxiliary complexing agent (c) calculate the molar concentration of the fe3+ in the original sample (d) calculate the molar concentration of the cu2+ in the original sample
A 25.00 mL aliquot of a solution containing Cu2+ and Fe3+ was titrated with 17.08 mL of 0.05095 M EDTA. A second 25.00 mL aliquot of the Cu/Fe mixture was treated with NaF to form a stable iron-fluoride complex. This mixture was then titrated with EDTA and the endpoint volume was found to be 5.47 mL. Calculate the amounts of Cu2+ and Fe3+ in mg/L. Molar mass (g/mol): Fe = 55.85 and Cu = 63.55
b. EDTA cannot be used as a primary standard always. When can EDTA be used as a primarystandard ? Give reason why EDTA is used in complexometric titrations.
A 25 ml solution of unkown containing Fe+3 and Ca+2 required 16.06 ml of 0.051 M EDTA for complete the reaction. A 50 ml of the same of the unkown was treated with NH4F to protect Fe+3. then the Ca+2 was reduced and masked by addition thiourea. upon addition of 25 ml of 0.051 M EDTA, the Fe+3 was liberated fron its floride complex and formed EDTA complex . the excess ETDA required 19.77 ml of 0.01883 M pb+2 to reach the end point by using xylenol orange indicator. Find the concentarion of Ca+2 in the unkown sample?
A 25.00 mL sample containing Fe3+ was treated with 10.00 mL of 0.03676 M EDTA to complex all the Fe3+ and leave excess EDTA in solution. The excess EDTA was then back-titrated, requiring 2.37 mL of 0.04615 M Mg2+. What was the concentration of Fe3+ in the original solution in ppm Fe3+?
Chromium (III) ion was complexed with EDTA and determined via back-titration. Given the following data, compute for %CrCl3(MM=158.35) in a 2.63 g sample. Conc. of EDTA titrant: 0.0103 MVolume used: 5.00 mLConc. of Zn2+ back titrant: 0.0112 MVolume used: 1.32 mL
What conditions must be followed to perform complexometric titration?A. presence of a catalystB. low temperatureC. [metal-EDTA] complex should be more stable than [metal-indicator] complexD. always add an EDTA solution to the analyzed solution and not vice versa
Aluminium (III) reacts with the Chelating agent HL to give a complex AlL3 that is readily soluble in CHCl3. A spectrophotometric study shows that when a 1.77 x10-4 M aqueous solution of aluminium was extracted with CHCl3 containing 0.01M HL, the analytical concentration of aluminium at pH 4.5 in the two phases were 1.5 x 10-4 M [Al3+]upper) and 2.73 x10-5 M ([AlL3]lower) . What is the distribution coefficient at pH 4.8.
A 50.0-mL sample containing Ni2+ was treated with 25.0 mL of 0.0500 M EDTA to complex all the Ni2+ and leave excess EDTA in solution. The excess EDTA was then back-titrated, requiring 5.00 mL of 0.0500 M Zn2+. What was the molar concentration of Ni2+ in the original solution?
From the structure of EDTA (and knowing that it is ethylene diamine tetraacetic acid) explain why the EDTA ismuch more soluble after the addition of the basic ammonia buffer. Is there a chemical reaction that isresponsible for this solubility change? If so explain what the reaction is and how it would be responsiblefor the increased solubility of EDTA in water.
A 25.00-mL sample containing Fe3+ was treated with 10.00 mL of 0.036 7 M EDTA to complex all the Fe3+ and leave excess EDTA in solution. The excess EDTA was then back titrated, requiring 2.37 mL of 0.0461 M Mg2+. What was the concentration of Fe3+ in the original solution?
The red color of soil is often due to the presence of iron. Metal ions are extracted from soil by stirring the soil in acid and then filtering the solution. One method for the analysis of Fe2+ is to form the highly colored Fe2+–thioglycolic acid complex. The complex absorbs strongly at 535 nm. Calibration standards of 1.00, 2.00, 3.00, 4.00, and 5.00 ppm are prepared by transferring appropriate amounts of a 10.0 ppm working solution of Fe2+ into separate 50-mL volumetric flasks, each of which contains 5 mL of thioglycolic acid, 2 mL of 20% w/v ammonium citrate, and 5 mL of 0.22 M NH3. After diluting to volume and mixing, the absorbances of the standards are measured. a)What is the effect on the reported concentration of iron in the sample if there is a trace impurity of Fe2+in the ammonium citrate?