26 Precipitation (Argentometric) Titrations

epgp books

 

1. What are precipitation reactions?

2. What are solubility product and ionic product?

3. What are the factors effecting solubility?

4. What are argentometric titrations?

5. How is the end point of precipitation titration?

6. What are the applications of precipitation titration?

 

1. Precipitation Titration: Precipitation titrations involve the stoichiometric reaction of standard precipitating agent with analyte. Precipitation titrations involving the use of silver ion as precipitating agent are classified as argentometric titrations.

 

1 Basic Requirements: The precipitate formed should have a definite stoichiometric.

• The equilibrium between the precipitate and its ions in solution much be attained rapidly.

• The precipitate must be of low solubility in the solution i.e., low solubility product.

•  Suitable method for the detection of stoichiometric end-point of the titration.

 

2 Solubility Product:

What happens when a sparingly soluble salt is put into a solvent?

Figure 1: Stages of dissolution of salt

An equilibrium exists between a partially soluble substance and its solution

For Example
BaSO4 (s)⇌ Ba2+ (aq) + SO42- (aq)

K  = [Ba2+ (aq)][SO42- (aq)] / [BaSO4(s)] K = Ksp, the solubility-product constant, as [BaSO4 (s)] = unity

Ksp = [Ba2+(aq)][SO42-(aq)]

Thus, Solubility Product (Ksp) is defined as equilibrium constant for the equilibrium between an ionic solid and its saturated solution.

 

2.1 General Expression for Solubility Product:

AaBb(s) aAb+ (aq) + bBa- (aq)

Ksp = [Ab+]a [Ba-]b

Example: PbI2 (s)  Pb2+ + 2 I-Ksp = [Pb2+] [I-]2

Higher Ksp indicates high molar solubility; whereas lower Ksp indicates low molar solubility of the salt e.g. AgCl has lower solubility than NaCl.

 

3 Conditions for Precipitation:

  •  Precipitation is reverse of the dissolution phenomenon

 

3.1 Ionic Product: ionic product is the term given for the product of the concentration of cation and anion raised to the power of their respective stoichiometry in the salt.

  • Ionic product is the evaluation of the strength of ions under non-equilibrium condition. e.g. Ionic Product of BaCl2 would be given by Qsp = [Ba2+(aq)][Cl-(aq)]2
  • Qsp = Ksp è saturated solution, but no precipitate
  • Qsp > Ksp è saturated solution, with precipitate
  • Qsp < Ksp è unsaturated solution,

 

3.2 Factors Affecting Solubility / Precipitation

•  Temperature

•  Common ion effect

•  Salt effect

•  pH of solution

•  Formation of complex ion

3.2.1     Common Ion Effect

•  Consider the following solubility equilibrium:

AgCl(s) ⇌ Ag+(aq) + Cl-(aq); Ksp = 1.6 x 10-10

•  The solubility of AgCl is 1.3 x 10-5 mol/L at 25oC.

•  If NaCl is added, equilibrium shifts left due to increase in [Cl–] and more AgCl will precipitate out. For example, if [Cl-] = 1.0 x 10-2 M,

•  Solubility of AgCl = (1.6 x 10-10)/(1.0 x 10-2)=  1.6 x 10-8 mol/L

 

3.2.2    Effect of pH on Solubility

  • Consider the following equilibrium: Ag3PO4(s) ⇌ 3Ag+(aq) + PO43-(aq);
  •  If HNO3 is added, the following reaction occurs: H3O+(aq) + PO43-(aq) ⇌ HPO42-(aq) + H2O
  • Consider the following equilibrium: Mg(OH)2(s) ⇌ Mg2+(aq) + 2 OH-(aq);
  •  Increasing the pH means increasing [OH-] and equilibrium will shift to the left, causing some of Mg(OH)2 to precipitate out.

3.2.3    Formation of Complex Ions on Solubility

  • Many transition metals ions have strong affinity for ligands to form complex ions.
  • Ligands are molecules, such as H2O, NH3 and CO, or anions, such as F-, CN- and S2O32-.
  • Complexes are soluble – thus, the formation of complex ions increases solubility of slightly soluble ionic compounds.
  • Consider the following equilibria:
  • Combining the two equations yields:AgCl(s) + 2NH3(aq) ⇌ Ag(NH3)2+(aq) + Cl-(aq); Knet = Ksp x Kf = (1.6 x 10-10) x (1.7 x 107)=  2.7 x 10-3
  • Knet > Ksp implies that AgCl is more soluble in aqueous NH3 than in water.
  • Consider the following case: 20.0 mL of 0.025 M Pb(NO3)2 is added to 30.0 mL of 0.10 M NaCl. Predict if precipitate of PbCl2 will form. (Ksp for PbCl2 = 1.6 x 10-5)

Calculation:

[Pb2+] = (20.0 mL x 0.025 M)/(50.0 mL) = 0.010 M

[Cl–] = (30.0 mL x 0.10 M)/(50.0 mL) = 0.060 M

Qsp = [Pb2+][Cl–]2 = (0.010 M)(0.060 M)2  = 3.6 x 10-5

Qsp > Ksp  è precipitate of PbCl2 will form.

4 Argentometric Titration: Argentometric titration is carried out by addition of standardized AgNO3 solution to the analyte solution.Upon increasing the concentration of AgNO3 in the titre to an extent that the ionic product of Ag+ and Cl- exceeds the solubility product, a suspension is formed. This increase in the titrant leads to depletion of titre and increase in the –log10 [Cl-]. The titration curve is normally broken down in three regions for the purposes of calculation and a function for pCl is determined for each region:

4.1 Pre-equivalence region

4.2 Equivalence point

4.3 Post-equivalence region

Figure 2:     The precipitation titration response curve for argentometric titration of chloride ions Argentometric Titration:

 

Table 1: The progress of argentometric titration for chloride ions Calculation at S. No. 2: 10.00 mL Titrant Added (Pre-equivalence Region)

mmol Cl– started = 25.00 mL 0.100 M = 2.50 mmol

 

5  Precipitation Titration: End Point Detection: End point determination in a precipitation titration could be done by four methods

–  Mohr’s Method

–  Volhards Method

–  Fajan’s Method

–  Conductivity Method

End Point Indication:

Two types of indicators are used to determine the end point in precipitation titrations. One type of indicators are those which form colored compounds when excess amount of titrant is added. Adsorption indicators are the indicators which are adsorbed on the precipitates at equivalence point with change in color on adsorption.

End point detection with indicators that forms colored compounds with excess titrants

5.1 Mohr’s Method:

Direct precipitation titration e.g. determination of chloride using silver nitrate in presence of chromate as indicator.

NaCl + AgNO3 ⇆ AgCl (s) + Na+ (aq.) + NO3– (aq) Ksp = 1.82 x 10-10

Chromate does not precipitate until Cl- ions have been completely consumed by Ag+ ions owing to solubility product and concentration of chromate in solution. The pH of the solution plays an important role with chromate under acidic conditions undergo a change to dichromate, whereas as pH above 10 silver precipitates as AgOH.

5.2 Volhard’s Method: Indirect Argentometric titration method for halides utilizes the reaction of thiocyanate with Ag+ ions to form precipitate. These titrations are performed in dilute nitric acid as all the silver halides are sparingly soluble in acidic solution. A measured amount of AgNO3 solutions is added to the sample

Excess of silver ions determined by back titration with thiocyanate ions in presence of Fe(III) ions as indicator

For titrations where AgX (I, Br) is less soluble than AgSCN there is no need to remove precipitates before back titration. For AgCl which is more soluble than AgSCN, the AgCl slowly dissolves and is replaced by AgSCN before the reaction of Fe+3 in solution. For this, more amount of titre is required resulting into erroneous result. These errors can be removed by addition of masking agent such as dibutylphthalate.

Adsorption indicators:

5.3 FAJAN’s METHOD:

The adsorption indicators are usually anionic dyes which get attached to the positively charged particles that are produced during the titration. The method exploits the surface charge on the colloidal precipitates formed. The dye is adsorbed on the surface of colloid based on the charge on the surface e.g. argentometric titration for estimation of chloride ions.

  • During the pre-equivalence stage, Cl- ions are more and the crystals of AgCl which ppt. have negative surface charge due the excess of chloride ions.
  • Upon further titration with Ag+ in the post equivalence stage, the crystals of AgCl have positive charge due to excess of Ag+ ions.
  • The dye is not adsorbed in the pre equivalence stage thus the solution has yellow color and precipitates are white. However in the post equivalence point the positively colloidal precipitates adsorbs fluoroscein with characteristics red color imparted to the precipitates.

Figure 3: Depicting the mechanism of action of Fajan’s indicator for argentometric titrations of chloride ions

5.4 Conductometric Argentometric Titration: As precipitation reactions involve the removal of ions from solution as salt precipitation leads to lowering of the conductivity of the titre initially.

  • In pre equivalence stage, chloride gets precipitated as AgCl and the conductivity of the sample decreases
  • At equivalence point, the chloride ions conductivity of the solution is minimum and is replaced by the bulky nitrate ions.
  • In post equivalence stage, excess of AgNO3 in solution increases the conductivity of the sample again.

Figure 4: The conductometric behavior of argentometric titration of chloride ions

6  Application of Precipitation Titration:

6.1 Determination of chloride (mg/L) in water and wastewater samples at low TDS. To a 100 ml water sample to be analyzed in the titre flask, add 5-6 drops of chromate indicator solution. Titrate with 0.0141 N silver nitrate solution to reddish brown end point. Note down the volume of titrant used.

6.2 Determination of cyanide (mg/L) in wastewater from acrylonitrile polymer production unit wastewater. CN- containing wastewater is mixed with KI under slightly basic conditions and argentometric titration performed.

– Initially, silver complexes to give silver cyanide complexes Ag+ + CN- à Ag(CN)2-

–  Excess Ag+ ions react with Iodide forming precipitate

6.3 Determination of sulphate in water and wastewater, upon precipitation as BaSO4 using conductometric titration

– Sulphate ions in a water/wastewater could be determined by a conductometric titration with BaCl2 solution.

– In pre equivalence stage, the sulphate gets precipitated and the conductivity of the sample decreases

–  At equivalence point, the sulphate ions conductivity of the solution is minimum.

– In post equivalence stage, excess of BaCl2 in solution increases the conductivity of the sample

Figure 5: The conductometric behavior of argentometric titration for estimation of suphate in water

you can view video on Precipitation(Argentometric) Titrations

Bibliography

  1. G. Marr and B.W. Rocket, ‘Practical Inorganic Chemistry’, University Science Books, 1999.
  2. G. Pass and H. Sutcliffe, ‘Practical Inorganic Chemistry’, Chapman and Hall, London, 1968.
  3. Vogel’s Textbook of Quantitative Chemical Analysis, Arthur Israel Vogel, Prentice Hall, 2000.
  4. J. Mendham, R. C. Denney, J. D. Barnes, M. Thomas, ‘Vogel’s Textbook of Quantitative Analysis’, Pearson Education, 2006.