D. Harvey - Modern Analytical Chemistry (794078), страница 74
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A temperatureof 110 °C is usually sufficient when removing water and other easily volatilized impurities. A conventional laboratory oven is sufficient for this purpose. Higher temperatures require the use of a muffle furnace, or a Bunsen or Meker burner, and arenecessary when the precipitate must be thermally decomposed before weighing orwhen using filter paper. To ensure that drying is complete the precipitate is repeatedly dried and weighed until a constant weight is obtained.Filter paper’s ability to absorb moisture makes its removal necessary beforeweighing the precipitate.
This is accomplished by folding the filter paper over theprecipitate and transferring both the filter paper and the precipitate to a porcelainor platinum crucible. Gentle heating is used to first dry and then to char the filterpaper. Once the paper begins to char, the temperature is slowly increased.
Althoughthe paper will often show traces of smoke, it is not allowed to catch fire as any precipitate retained by soot particles will be lost. After the paper is completely charredthe temperature is slowly raised to a higher temperature. At this stage any carbonleft after charring is oxidized to CO2.Fritted glass crucibles cannot withstand high temperatures and, therefore,should only be dried in an oven at temperatures below 200 °C. The glass fiber matsused in Gooch crucibles can be heated to a maximum temperature of approximately 500 °C.Composition of Final Precipitate The quantitative application of precipitationgravimetry, which is based on a conservation of mass, requires that the final precipitate have a well-defined composition.
Precipitates containing volatile ions or substantial amounts of hydrated water are usually dried at a temperature that is sufficient to completely remove the volatile species. For example, one standardgravimetric method for the determination of magnesium involves the precipitationof MgNH4PO4 6H2O. Unfortunately, this precipitate is difficult to dry at lowertemperatures without losing an inconsistent amount of hydrated water and ammonia. Instead, the precipitate is dried at temperatures above 1000 °C, where it decomposes to magnesium pyrophosphate, Mg2P2O7.⋅245peptizationThe reverse of coagulation in which acoagulated precipitate reverts to smallerparticles.1400-CH08 9/9/99 2:17 PM Page 246Modern Analytical ChemistryAn additional problem is encountered when the isolated solid is nonstoichiometric.
For example, precipitating Mn2+ as Mn(OH)2, followed by heatingto produce the oxide, frequently produces a solid with a stoichiometry of MnOx ,where x varies between 1 and 2. In this case the nonstoichiometric product resultsfrom the formation of a mixture of several oxides that differ in the oxidation stateof manganese. Other nonstoichiometric compounds form as a result of lattice defects in the crystal structure.6Representative Method The best way to appreciate the importance of the theoretical and practical details discussed in the previous section is to carefully examine theprocedure for a typical precipitation gravimetric method. Although each methodhas its own unique considerations, the determination of Mg2+ in water and wastewater by precipitating MgNH4PO4 6H2O and isolating Mg2P2O7 provides an instructive example of a typical procedure.⋅Representative Methods246Method 8.1Determination of Mg2+ in Water and Wastewater7⋅Description of Method.
Magnesium is precipitated as MgNH4PO4 6H2O using(NH4)2HPO4 as the precipitant. The precipitate’s solubility in neutral solutions(0.0065 g/100 mL in pure water at 10 °C) is relatively high, but it is much less solublein the presence of dilute ammonia (0.0003 g/100 mL in 0.6 M NH3). The precipitant isnot very selective, so a preliminary separation of Mg2+ from potential interferents isnecessary. Calcium, which is the most significant interferent, is usually removed byits prior precipitation as the oxalate. The presence of excess ammonium salts fromthe precipitant or the addition of too much ammonia can lead to the formation ofMg(NH4)4(PO4)2, which is subsequently isolated as Mg(PO3)2 after drying. Theprecipitate is isolated by filtration using a rinse solution of dilute ammonia.
Afterfiltering, the precipitate is converted to Mg2P2O7 and weighed.Procedure. Transfer a sample containing no more than 60 mg of Mg2+ into a600-mL beaker. Add 2–3 drops of methyl red indicator, and, if necessary, adjust thevolume to 150 mL. Acidify the solution with 6 M HCl, and add 10 mL of 30% w/v(NH4)2HPO4. After cooling, add concentrated NH3 dropwise, and while constantlystirring, until the methyl red indicator turns yellow (pH > 6.3).
After stirring for5 min, add 5 mL of concentrated NH3, and continue stirring for an additional 10 min.Allow the resulting solution and precipitate to stand overnight. Isolate theprecipitate by filtration, rinsing with 5% v/v NH3. Dissolve the precipitate in 50 mLof 10% v/v HCl, and precipitate a second time following the same procedure. Afterfiltering, carefully remove the filter paper by charring.
Heat the precipitate at 500 °Cuntil the residue is white, and then bring the precipitate to constant weight at1100 °C.Questions1. Why does the procedure call for a sample containing no more than 60 mg ofMg2+?A sample containing 60 mg of Mg2+ will generate approximately 600 mg, or0.6 g, of MgNH4PO4 6H2O. This is a substantial amount of precipitate to workwith during the filtration step. Large quantities of precipitate may be difficultto filter and difficult to adequately rinse free of impurities.⋅—Continued1400-CH08 9/9/99 2:17 PM Page 247Chapter 8 Gravimetric Methods of Analysis2. Why is the solution acidified with HCl before the precipitant is added?⋅The HCl is added to ensure that MgNH4PO4 6H2O does not precipitate whenthe precipitant is initially added. Because PO43– is a weak base, the precipitate issoluble in a strongly acidic solution.
If the precipitant is added under neutral orbasic conditions (high RSS), the resulting precipitate will consist of smaller, lesspure particles. Increasing the pH by adding base allows the precipitate ofMgNH4PO4 6H2O to form under more favorable (low RSS) conditions.⋅3. Why is the acid–base indicator methyl red added to the solution?The indicator’s color change, which occurs at a pH of approximately 6.3,indicates when sufficient NH3 has been added to neutralize the HCl added atthe beginning of the procedure. The amount of NH3 added is crucial to thisprocedure.
If insufficient NH3 is added, the precipitate’s solubility increases,leading to a negative determinate error. If too much NH3 is added, theprecipitate may contain traces of Mg(NH4)4(PO4)2, which, on ignition, formsMg(PO3)2. This increases the mass of the ignited precipitate, giving a positivedeterminate error.
Once enough NH3 has been added to neutralize the HCl,additional NH3 is added to quantitatively precipitate MgNH4PO4 6H2O.⋅4. Explain why the formation of Mg(PO3)2 in place of Mg2P2O7 increases the massof precipitate.The desired final precipitate, Mg2P2O7, contains two moles of Mg, and theimpurity, Mg(PO3)2, contains only one mole of Mg. Conservation of mass,therefore, requires that two moles of Mg(PO3)2 must form in place of each moleof Mg2P2O7.
One mole of Mg2P2O7 weights 222.6 g. Two moles of Mg(PO3)2weigh 364.5 g. Any replacement of Mg2P2O7 with Mg(PO3)2 must increase theprecipitate’s mass.5. What additional steps in the procedure, beyond those discussed in questions 2and 3, are taken to improve the precipitate’s purity?Two additional steps in the procedure help form a precipitate that is free ofimpurities: digestion and reprecipitation.6. Why is the precipitate rinsed with a solution of 5% v/v NH3?This is done for the same reason that precipitation is carried out in anammonical solution; using dilute ammonia minimizes solubility losses whenrinsing the precipitate.8B.2 Quantitative ApplicationsAlthough not in common use, precipitation gravimetry still provides a reliablemeans for assessing the accuracy of other methods of analysis or for verifying thecomposition of standard reference materials.
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