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Precipitation Titration

Precipitation Titration  Analyte is titrated with a standard solution of a precipitating agent in accordance with defined reaction stoichiometry.

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Precipitation Titration

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  1. Precipitation Titration  Analyte is titrated with a standard solution of a precipitating agent in accordance with defined reaction stoichiometry.  Detection of the end point (at completion of the precipitation) is usually by either the appearance of excess titrant or the disappearance of the reactant. Formation of insoluble metal-ligand precipitates. ·Argentometric Methods Titration methods based upon utilizing silver nitrate as a precipitating agent. ·Silver ion is extremely useful in precipitation reactions including: -Halides (Cl, Br, I) -Pseudohalides (S2, HS,CN, SCN)

  2. Requirements 1.The precipitate formation is stoichiometric.  2.To allow the titrant to be added quickly, the equilibrium between the precipitate and its ions in solution much be attained rapidly. A slow attainment of equilibrium will cause overtitration.  3.The precipitate must be of low solubility in the solution. This is indicated by small equilibrium constant (Ksp).  4.A method to detect the stoichiometric point of the titration must be available. Although a number of indicators are available, in general, the best method for detecting the end point in precipitation titration is by an instrumental technique.

  3. Titration Curves for Argentometric Methods ·Plots of titration curves are normally sigmoidal curves consisting of pAg (or pAnalyte) versus volume of AgNO3 solution added. Example: Titration of chloride with silver. The points on the curve can be calculated, given the analyte concentration, AgNO3 concentration and the appropriate Ksp.

  4. Titrant Added (mL) pAg 0.00 1.00 10.00 1.48 20.00 4.87 25.00 7.79 ·The titration curve is normally broken down in three regions for the purposes of calculation and a function for pAg is determined for each region: 1.Pre-equivalence region 2.Equivalence point 3.Post-equivalence region Example  Calculate the [Ag+], [Cl] and pAg in a titration of 20.00 mL of 0.100 M AgNO3 with 0.100 M NaCl.

  5. a) 0.00 mL Titrant Added (Pre-equivalence Region)  At the start of the titration, [Ag+] = 0.100 M. pAg = log [Ag+] = log 0.100 = 1  b) 10.00 mL Titrant Added (Pre-equivalence Region)  mmol Ag+ started = 20.00 mL  0.100 M = 2.00 mmol mmol Cl added = 10.00 mL  0.100 M = 1.00 mmol  mmol Ag+ left = 2.00  1.000 = 1.00 mmol Volume of solution = 20 + 10 = 30 mL  [Ag+] = (1.00 mmol)/30 mL = 0.0333 pAg = - log 0.0333 = 1.48

  6. c) 20.00 mL Titrant Added (Equivalence Point)  mmol Ag+ started = 20.00 mL  0.100 M = 2.00 mmol mmol Cl added = 20.00 mL  0.100 M = 2.00 mmol  [Ag+] = [Cl] = S Ksp = [Ag+][Cl ] 1.8  1010 (M)2 = S  S S = 1.34  105 M  pAg = pCl =  log 1.34  105 = 4.87 d) 25.00 mL Titrant Added (Post-equivalence Region)  After the stoichiometric point, excess NaCl is being added into the solution. After 25 mL of 0.100 M NaCl has been added,

  7. mmol Ag+ started = 20.00 mL  0.100 M = 2.00 mmol mmol Cl added = 25.00 mL  0.100 M = 2.50 mmol  mmol Cl in excess = 2.50  2.000 = 0.50 mmol Volume of solution = 20 + 25 = 45 mL [Cl]  (0.50 mmol)/45 mL = 0.0111* Ksp = [Ag+][Cl ] 1.8  1010 (M)2 = S  [0.0111] S =1.62  108 M pAg = - log 1.62  108 = 7.79  * Note that in this calculation, [Cl] contributed by the solubility of the AgCl is considered negligible. This is a reasonable approximation as S << [Cl]. A more accurate expression for the chloride concentration is [Cl] = 0.0111 + S

  8. Observations About Argentometric Titrations ·High reagent concentrations give sharper, more dramatic equivalence point changes in pAg and better endpoints. The smaller the Ksp, the more complete the precipitation reaction and the sharper the equivalence region changes.  Both Ksp and the reagent concentrations affect the choice and use of an endpoint indicator. Endpoint Indicators for Argentometric Titrations ·Indicators for argentometric titrations are selected to produce a color change at or near the equivalence point. ·Normally the indicator is selected to react with the added titrating agent, not the analyte. If A is the analyte, R the titrating agent and In is the indicator. A + R AR(s) In(color) + R InR(new color) To make the indicator change color, excess R must be added. Obviously, the smaller the excess added to cause the color change, the smaller the endpoint error. meThis means that the indicator should give large color changes at very low concentrations.

  9. End Point Determination in Precipitation Titrations  Titrations with Ag+ are called argentometric titrations.  For argentometric titrations, three classical methods based on color indicators can be used for end point detection: ·Mohr titration – formation of colored precipitate at the end point. ·Volhard titration – formation of a soluble, colored complex at the end point. ·Fajans titration – adsorption of a colored indicator on the precipitate at the end point.

  10. Mohr method The Mohr method was first published in 1855 as a method for chloride analysis. ·In the precipitation of chloride by silver ion, chromate ion (CrO42) is used as an indicator in the formation of Ag2CrO4, a reddish-brown precipitate formed when excess Ag+ is present.  Ag+ + Cl AgCl(s) titrant analyte white precipitate Ksp= 1.8 x 10-10 (S = 1.34 x 10-5 M) ·Mohr indicator reaction (end point), 2Ag+ + CrO42 Ag2CrO4(s) titrant indicator red precipitate Ksp= 1.2 x 10-12 (s = 6.7 x 10-5 M)

  11. ·The concentration of titrant rises sharply near the equivalence point, and the solubility of Ag2CrO4 is exceeded. ·The titrations are performed only in neutral or slightly basic medium to prevent silver hydroxide formation (at pH > 10). 2Ag+ + 2OH 2AgOH(s) Ag2O(s) + H2O precipitate ·Or the formation of chromic acid at pH < 7.  CrO42 + H3O+ HCrO4 + H2O 2 CrO42 + 2 H3O+ Cr2O72 + H2O [CrO42] become lower, more Ag+ to be added to reach end point, which cause error. Suitable pH??

  12. Volhard METHOD    Determination of Cl. For titration of silver ion with thiocyanate (SCN) and iron(III) as an indicator.First published in 1874. Reactions: Ag+ + Cl AgCl(s) Ksp = 1.82 x 10-10 titrant #1 white precipitate (excess) Ag+ + SCN AgSCN(s) Ksp = 1.1 x 10-12 titrant #2 white precipitate Fe3+ + SCN FeSCN2+ Kf = 1.4 x 10+2 indicator red precipitate

  13. ·The endpoint is routinely used for halide determinations where a known excess of silver ion is added to precipitate the halide ion, and the excess silver ion is back titrated using the thiocyanate/iron(III) as the indicator. ·The silver chloride precipitate is filtered, and the excess silver ion is titrated with thiocyanate producing a white precipitate of AgSCN. ·Once the silver is consumed, the excess thiocyante reacts with the iron(III) ion producing a red FeSCN2+ complex. Thus, the appearance of the red color at the endpoint. The red FeSCN2+ complex color is detectable at 6.4 x 10-6 M concentrations and above.  The titration is usually done in acidic pH medium to prevent precipitation of iron hydroxides, Fe(OH)3.

  14. ·The AgCl precipitate must be separated from the thiocyanate to prevent the reaction: AgCl + SCN- AgSCN + Cl- ·Since AgSCN is less soluble than AgCl, equilibrium will shift to the right causing a negative error for the chloride analysis. To eliminate this error, AgCl must be filtered or add nitrobenzene before titrating with thiocyanate; nitrobenzene will form an oily layer on the surface of the AgCl precipitate, thus preventing its reaction with thiocyanate.

  15. Fajans Method (Adsorption Indicators) ·Adsorption indicators are organic compounds that tend to be adsorbed onto the surface of the solid precipitate in a precipitation titration. ·Adsorption indicators work best when: -They do not precipitate out silver ion when the indicators are at low concentration. -They bind to the precipitate only when excess silver ion is present to produce color.  Example of Adsorption Indicators: Fluorescein ·A polycyclic compound that ionizes in solution to yield yellow-green fluoresceinate ions.  Fluoresceinate adsorbs to silver ions on the surface of a precipitate when excess silver ion is present, producing a reddish-colored surface.

  16. ·Only the ionized fluoresceinate produces the red color. ·Titration of NaCl with AgNO3   Ag+ + Cl AgCl(s) titrant analyte white precipitate ·During the titration, colloids are formed. ·Before the equivalence point, the surface of the precipitant particles will be negatively charged due to the adsorption of excess Cl the surface of the particles. A diffuse positive counter-ion layer will surround the particles. ·The primary adsorption layer is negatively charged and the anionic indicator is repelled.

  17. ·Before the end point (Cl in excess). (AgCl) • Cl M+ 1 Layer 2 Layer When the equivalence point is reached, there is no longer an excess of analyte Cl, and the surface of the colloidal particles are largely neutral.

  18. ·After the equivalence point, there will be an excess of titrant Ag+, some of these will adsorb to the AgCl particles, which will now be surrounded by a diffuse negative counterion layer. ·Now, negatively charged fluorescein can penetrate the counter ion layer and adsorb onto the AgCl lattice due to its affinity to Ag+.

  19. (AgCl) • Ag+ X  1 Layer 2 Layer After the end point (Ag+ in excess). ·The fluorescein concentration is too small to form an AgIn precipitate in solution. The formed complex has a red color. ·Conditions that favor the formation of colloids: -Fast addition of reagents. -Add dextrin to retard coagulation. -   pH control is important to generate the dye anion in solution. -Avoid light, since Ag-dye complex is light sensitive and decomposes after a while.

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