Precipitation Titrimetry
Precipitation titrimetry is based upon reactions that yield ionic compounds of limited solubility. The slow rate of precipitate formation limits the number of precipitating agents that can be used in titrimetry to a handful.
A reaction in which the analyte and titrant form an insoluble precipitate also can form the basis for a titration. We call this type of titration a precipitation titration.
One of the earliest precipitation titrations, developed at the end of the eighteenth century, was for the analysis of K2CO3 and K2SO4 in potash. Calcium nitrate, Ca(NO3)2, was used as a titrant, forming a precipitate of CaCO3 and CaSO4. The end point was signalled by noting when the addition of titrant ceased to generate additional precipitate. The importance of precipitation titrimetry as an analytical method reached its zenith in the nineteenth century when several methods were developed for determining Ag+ and halide ions.
Requirements to reactions and defined substances:
§ The defined substance should be dissolved in water and give an ion which would be active in sedimentation reaction.
§ The received precipitate should be practically insoluble (Ksp<10-8 ¸ – 10, S<10-5 mol/L).
§ Results of titration should not be deformed by the adsorption phenomena (coprecipitation).
§ Precipitate should form enough quickly.
§ There should be a possibility of fixing of an equivalence point.
Classification of methods precipitation titration (on titrant):
1. Argentometry
2. Thiocyanatometry
3. Mercurometry
4. Sulphatometry
5. Hexacianoferratometry
Titration Curves
Titration curves for a single anion are derived in a way completely analogous to another titration methods. The only difference is that the solubility product of the precipitate is substituted to for the ion-product constant for water.
The change in p-function value at the equivalence point becomes grater as the solubility products become smaller – that is, as the reaction between the analyte and precipitant becomes more complete. Ions forming precipitates with solubility products much larger than about 10-10 do not yield satisfactory end point.
The titration curve for a precipitation titration follows the change in either the analyte’s or titrant’s concentration as a function of the volume of titrant. For example, in an analysis for I– using Ag+ as a titrant
Ag+ + I– ® AgI¯
the titration curve may be a plot of pAg or pI as a function of the titrant’s volume.
As we have done with previous titrations, we first show how to calculate the titration curve and then demonstrate how to quickly sketch the titration curve.
Calculating the Titration Curve. As an example, let’s calculate the titration curve for the titration of 50.0 mL of 0.0500 M Cl– with 0.100 M Ag+. The reaction in this case is
Ag+ + Cl– ® AgCl¯
The equilibrium constant for the reaction is
K = (Ksp)–1 = (1.8 ´10–10)–1 = 5.6 ´109
Since the equilibrium constant is large, we may assume that Ag+ and Cl– react completely.
By now you are familiar with our approach to calculating titration curves. The first task is to calculate the volume of Ag+ needed to reach the equivalence point. The stoichiometry of the reaction requires that
Moles Ag+ = Moles Cl–
or
MAgVAg = MClVCl
Solving for the volume of Ag+
shows that we need 25.0 mL of Ag+ to reach the equivalence point.
Before the equivalence point Cl– is in excess. The concentration of unreacted Cl– after adding 10.0 mL of Ag+, for example, is
If the titration curve follows the change in concentration for Cl–, then we calculate pCl as
pCl = –log[Cl–] = –log(2.50 ´10–2) = 1.60
However, if we wish to follow the change in concentration for Ag+ then we must first calculate its concentration. To do so we use the Ksp expression for AgCl
Ksp = [Ag+][Cl–] = 1.8 ´10–10
Solving for the concentration of Ag+
gives a pAg of 8.14.
At the equivalence point, we know that the concentrations of Ag+ and Cl– are equal. Using the solubility product expression
Ksp = [Ag+][Cl–] = [Ag+]2 = 1.8 ´10–10
gives
[Ag+] = [Cl–] = 1.3 ´10–5 M.
At the equivalence point, therefore, pAg and pCl are both 4.89.
After the equivalence point, the titration mixture contains excess Ag+. The concentration of Ag+ after adding 35.0 mL of titrant is
or a pAg of 1.93. The concentration of Cl– is
or a pCl of 7.82.
Additional results for the titration curve are shown in Table and Figure.
Data for Titration of 50.0 mL of 0.0500 M Cl– with 0.100 M Ag
|
Volume AgNO3(mL) |
pCl |
pAg |
|
0.00 |
1.30 |
— |
|
5.00 |
1.44 |
8.31 |
|
10.00 |
1.60 |
8.14 |
|
15.00 |
1.81 |
7.93 |
|
20.00 |
2.15 |
7.60 |
|
25.00 |
4.89 |
4.89 |
|
30.00 |
7.54 |
2.20 |
|
35.00 |
7.82 |
1.93 |
|
40.00 |
7.97 |
1.78 |
|
45.00 |
8.07 |
1.68 |
|
50.00 |
8.14 |
1.60 |
|
The factors which define value of inflection points of titration on curves of precipitation titration
§ Concentration of titrant solutions and a defined ion (than more concentration, the titration inflection point is more)
§ Solubility of a precipitate (than solubility less, the titration inflection point is more)
§ Temperature (than more temperature, the solubility of a precipitate will be more and the inflection point is less)
§ Ionic strength of a solution (than more ionic strength of a solution, the solubility of a precipitate will be more and the inflection point is less)
Selecting and Evaluating the End Point
Initial attempts at developing precipitation titration methods were limited by a poor end point signal. Finding the end point by looking for the first addition of titrant that does not yield additional precipitate is cumbersome at best. Two types of end points are encountered in titration with precipitants:
1) chemical,
2) instrumental: potentiometric and amperometric.
Chemical indicators for precipitation titration
The end point produced by a chemical indicator usually consists of a colour change or, occasionally, the appearance or disappearance of turbidity in the solution being titrated. The requirements for an indicator for a precipitation titration are analogous to those for an indicator for a neutralisation titration:
1) the colour change should occur over the reagent limited range in p-function of the reagent or the analyte and
2) the colour change should take place within the steep portion of the titration curve for the analyte.
Argentometry
Argentometry used for halide-like anions (Hal–, CN–, SCN–) determination, which forms slightly soluble compounds with Ag+ ion. Standard titration solution is AgNO3. For back-titration uses standard solution of NaCl, that titrated AgNO3 surplus. The silver nitrate solution can be standardised against primary-standard-grade sodium chloride.
Four indicators have found extensive use for argentometric titration:
1. Chromate ion. The Mohr method.
The first important visual indicator to be developed was the Mohr method for Cl– using Ag+ as a titrant. By adding a small amount of K2CrO4 to the solution containing the analyte, the formation of a precipitate of reddish-brown Ag2CrO4 signals the end point. Because K2CrO4 imparts a yellow colour to the solution, obscuring the end point, the amount of CrO42– added is small enough that the end point is always later than the equivalence point. To compensate for this positive determinate error an analyte-free reagent blank is analyzed to determine the volume of titrant needed to effect a change in the indicator’s colour. The volume for the reagent blank is subsequently subtracted from the experimental end point to give the true end point. Because CrO42– is a weak base, the solution usually is maintained at a slightly alkaline pH. If the pH is too acidic, chromate is present as HCrO4–, and the Ag2CrO4 end point will be in significant error. The pH also must be kept below a level of 10 to avoid precipitating silver hydroxide.
Sodium chromate serves as indicator for the argentometric detrmination of chloride, bromide, and cyanide ions by reacting with silver cation to form a brick-red silver chromate (Ag2CrO4) precipitate in the equivalence point region:
AgNO3 + NaCl = AgCl¯ + NaNO3
white
AgNO3 + Na2CrO4 = Ag2CrO4¯ + NaNO3
red
The Mohr titration must be carried out at a pH between 7 and 10, because chromate ion is the conjugate base of the weak chromic acid. In more acidic solutions the chromate ion concentration is too low to produce the precipitate at the equivalence point. A suitable pH is achieved with sodium hydrogen carbonate.
The Mohr titration caot be used for iodide and thiocyanate determination, because these ions form colloid solutions with silver ion.
2. Iron(III) ion. The Volhard method.
A second end point is the Volhard method in which Ag+ is titrated with SCN– in the presence of Fe3+. Silver ions are titrated with a standard solution of potassium or ammonium thiocyanate:
AgNO3 + NaCl = AgCl¯ + NaNO3
white
AgNO3 + KSCN = AgSCN¯ + KNO3
white
Iron(III) ion serve as the indicator. The end point for the titration reaction
Ag+ + SCN– ® AgSCN¯
is the formation of the reddish coloured Fe(SCN)3 complex:
SCN– + Fe3+ ® Fe(SCN)3
The solution turns red with the first slight excess of thiocyanate ion:
2NH4Fe(SO4)2 + 6KSCN = 2Fe(SCN)3 + 3K2SO4 + (NH4)2SO4
red
The titration must be carried out in distinct acidic solution
1) to prevent precipitation of iron(III) as the hydrated oxide,
2) and such ions as carbonate, oxalate, and arsenate not interfere with silver ion.
The most important application of the Volhard method is for the indirect determination of halide ions. Sometime the indirect Volhard method called thiocyanometry. A measurement excess of standard silver nitrate solution is added to the sample, and the excess silver ion is determined by back-titration with a standard thiocyanate solution. At chloride determination occurs titration error, because silver chloride is more soluble than silver thiocyanate.
3. Adsorption indicators. The Fajans method.
A third end point is evaluated with Fajans’ method, which uses an adsorption indicator whose colour when adsorbed to the precipitate is different from that when it is in solution. The adsorption indicator is an organic compound that tends to be adsorbed onto the surface of the solid in a precipitation titration. The adsorption occurs near the equivalence point and colour transfers from the solution to the solid.
For example, when titrating Cl– with Ag+ the anionic dye dichlorofluoroscein is used as the indicator. Before the end point, the precipitate of AgCl has a negative surface charge due to the adsorption of excess Cl–. The anionic indicator is repelled by the precipitate and remains in solution where it has a greenish yellow colour. After the end point, the precipitate has a positive surface charge due to the adsorption of excess Ag+. The anionic indicator now adsorbs to the precipitate’s surface where its colour is pink. This change in colour signals the end point.
|
|
|
|
Tetrabromofluorescein (eosine) |
Fluoresceine |
Fluorescein is a typical adsorption indicator that is useful for the titration of chloride ion with silver nitrate. In aqueous solution, fluorescein partially dissociates into hydronium ions and negatively charged fluoresceinate ions that are yellow-green. The fluoresceinate ion forms an intensively red silver salt at equivalence point. For bromide, iodide, and thiocyanate titration as an indicator is used tetrabromfluorescein, named eosin.
Titration involving adsorption indicators are rapid, accurate, and reliable, but their application is limited to the relatively few precipitation reactions in which a colloidal precipitate is formed rapidly.
There are also so called the Gay-Lussac argentometric titration method without indicators. At equivalence point the titrated halide solution clears up because occurs coagulation of precipitate.
Mercurometry
For mercurometric titration is using mercury(I) salts for halide ion determination.
Hg2(NO3)2 + 2NaCl = Hg2Cl2¯+ 2NaNO3
white
Indicator is iron(III) thiocyanate, which disappearance at equivalence point:
3Hg2(NO3)2 + 2Fe(SCN)3 = 3Hg2(SCN)2¯+ 2Fe(NO3)3
red solution white
Also as an indicator can be used diphenylcarbazon, which is an adsorption indicator and at equivalence point change colour of precipitate from white to blue.
Superiority of mercurometry:
1. Not required expensive silver salts.
2. Mercury(I) precipitates are less soluble than analogous silver salts. Therefore, equivalence point is clearly marked.
3. Mercurometric determination are carried out with direct titration in acidic solution.
4. Chloride ion can be determined with reducers (S2–, SO32–) and oxidisers (MnO4–, Cr2O7–) presence.
5. Titration can be carried out in turbid and coloured solutions.
Thiocyanatometry
§ This is a precipitation titration in which SCN- is the titrant.
§ Titrant: ammonium or potassium thiocyanide NH4SCN, KSCN – secondary standard solution
§ Stardadization: on primary standard solution of AgNO3:
AgNO3 + NH4SCN = AgSCN¯ + NH4NO3
§ Indicator by standardization of ammonium or potassium thiocyanide with iron (ІІІ) salts:
Fe3+ + SCN- = [Fe(SCN)]2+
§ Medium: in presence of nitric acid
§ Indicator: iron (ІІІ) salts NH4Fe(SO4)2×12H2O in presence of nitric acid
Determinate substance: drugs, which contain Silver (Albumosesilber, colloid silver – Kollargol, silver nitrate).
!!! At the analysis of drugs which contaionionic silver, preliminary it is heated with sulphatic and nitric acids (receive ionic compound).
!!! At definition of iodides the indicator is added in the end of titration to avoid parallel:
2Fe3+ + 2I– = 2Fe2+ + I2
Another Precipitation Titration Methods
1. Barium determination with sulphate:
BaCl2 + H2SO4 = BaSO4¯ + 2HCl
Indicator is sodium rhodizonate, which disappearance red colour of solution.
2. Lead determination with chromate:
Pb(NO3)2 + K2CrO4 = PbCrO4¯ + 2KNO3
Indicator is AgNO3 solution. At equivalence point forms red precipitate.
3. Zinc determination with K4[Fe(CN)6]:
3ZnCl2 + K4[Fe(CN)6] = Zn3K2[Fe(CN)6] ¯ + 2KCl
Indicator is UO2(NO3)2, which forms brown precipitate with K4[Fe(CN)6].
Finding the End Point Potentiometrically
Another method for locating the end point of a precipitation titration is to monitor the change in concentration for the analyte or titrant using an ion-selective electrode. The end point can then be found from a visual inspection of the titration curve.
Determinate the end-point by potenthiometric way
Quantitative Applications
Precipitation titrimetry is rarely listed as a standard method of analysis, but may still be useful as a secondary analytical method for verifying results obtained by other methods. Most precipitation titrations involve Ag+ as either an analyte or titrant. Those titrations in which Ag+ is the titrant are called argentometric titrations. Table provides a list of several typical precipitation titrations.
Representative Examples of Precipitation Titrations
|
Analyte |
Titranta |
End Pointb |
|
AsO43– |
AgNO3, KSCN |
Volhard |
|
Br– |
AgNO3 |
Mohr or Fajans |
|
|
AgNO3, KSCN |
Volhard |
|
Cl– |
AgNO3 |
Mohr or Fajans |
|
|
AgNO3, KSCN |
Volhard* |
|
CO32– |
AgNO3, KSCN |
Volhard* |
|
C2O42– |
AgNO3, KSCN |
Volhard* |
|
CrO42– |
AgNO3, KSCN |
Volhard* |
|
I– |
AgNO3 |
Fajans |
|
|
AgNO3, KSCN |
Volhard |
|
PO43– |
AgNO3, KSCN |
Volhard* |
|
S2– |
AgNO3, KSCN |
Volhard* |
|
SCN– |
AgNO3, KSCN |
Volhard |
a) When two reagents are listed, the analysis is by a back titration. The first reagent is added in excess, and the second reagent is used to back titrate the excess.
b) For Volhard methods identified by an asterisk (*) the precipitated silver salt must be removed before carrying out the back titration.
Quantitative Calculations
The stoichiometry of a precipitation reaction is given by the conservation of charge between the titrant and analyte thus
