GRAVIMETRIC ANALYSIS.
Gravimetry based on measurement of weight of an analysed species or a compound containing the analysed species. Measuring mass is the most fundamental of all analytical measurements, and gravimetry is unquestionably the oldest analytical technique.
Gravimetric analysis describes a set of methods in analytical chemistry for the quantitative determination of an analyte based on the mass of a solid. A simple example is the measurement of solids suspended in a water sample: A known volume of water is filtered, and the collected solids are weighed.
Analytical balance
Advantages
Gravimetric analysis, if methods are followed carefully, provides for exceedingly precise analysis. In fact, gravimetric analysis was used to determine the atomic masses of many elements to six figure accuracy. Gravimetry provides very little room for instrumental error and does not require a series of standards for calculation of an unknown. Methods also do not require often expensive equipment. Gravimetric analysis, due to its high degree of accuracy, when performed correctly, can also be used to calibrate other instruments in lieu of reference standards.
Disadvantages
Gravimetric analysis usually only provides for the analysis of a single element, or a limited group of elements, at a time. Comparing modern dynamic flash combustion coupled with gas chromatography with traditional combustion analysis will show that the former is both faster and allows for simultaneous determination of multiple elements while traditional determination allowed only for the determination of carbon and hydrogen. Methods are often convoluted and a slight mis-step in a procedure can often mean disaster for the analysis (colloid formation in precipitation gravimetry, for example). Compare this with hardy methods such as spectrophotometry and one will find that analysis by these methods is much more efficient.
Gravimetric methods are divided into three groups: (1) mechanical separations; (2) dry methods; and (3) wet methods.
Mechanical Separations.—Under this head are classed the method of assaying tin ores, known as vanning, and the amalgamation assay for gold. A set of sieves to determine the relative proportion of powders of different degrees of fineness is sometimes useful. A set with 10, 20, 40 and 80 meshes to the inch is convenient.
Dry Assays.—An important distinction between wet and dry methods of assaying is, that in the former the substance is got into the liquid state by solution, whilst in the latter fusion is taken advantage of.
The difference between solution and fusion is easily illustrated: a lump of sugar heated over a candle-flame melts or fuses; suspended in water it dissolves. Many substances which are insoluble or infusible of themselves, become soluble or fusible when mixed with certain others; thus, in this way, solution is got with the aid of reagents, and fusion with the help of fluxes. For example, lead is insoluble in water, but if nitric acid be added, the metal rapidly disappears. It is convenient, but somewhat inaccurate, to say that the acid dissolves the lead. If the lead be acted on by nitric acid alone, without water, it is converted into a white powder, which does not dissolve until water is added; in this case it is obvious that the water is the solvent. The function of the acid is to convert the lead into a soluble compound.
Fluxes may act as true solvents. Fused carbonate of soda dissolves baric carbonate, and perhaps in many slags true solution occurs; but in the great majority of cases a flux is a solid reagent added for the purpose of forming a fusible compound with the earthy or stony minerals of the ore. Few of the minerals which occur in the gangue of an ore are fusible; and still fewer are sufficiently fusible for the purposes of the assayer, consequently the subject is one of importance, and it ought to be treated on chemical principles.
Using Mass as a Signal
There are two ways to use mass as an analytical signal. We can, of course, measure an analyte’s mass directly by placing it on a balance and recording its mass. For example, determination of the total suspended solids in water released from a sewage-treatment facility. Suspended solids are just that; solid matter that has yet to settle out of its solution matrix. The analysis is easy. A sample collects and passes it through a preweighed filter that retains the suspended solids. After drying to remove any residual moisture, the filter weighs. The difference between the filter’s original mass and final mass gives the mass of suspended solids. It is a direct analysis because the analyte itself is the object being weighed.
If the analyte is an aqueous ion, such as Pb2+, we cannot isolate the analyte by filtration because the Pb2+ is dissolved in the solution’s matrix. We can still measure the analyte’s mass, however, by chemically converting it to a solid form. If we suspend a pair of Pt electrodes in our solution and apply a sufficiently positive potential between them for a long enough time, we can force the reaction
Pb2+ + 4H2O « PbO2 + H2 + 2H3O+
to go to completion. The Pb2+ ion in solution oxidizes to PbO2 and deposits on the Pt electrode serving as the anode. If we weigh the Pt anode before and after applying the potential, the difference in the two measurements gives the mass of PbO2 and, from the reaction’s stoichiometry, the mass of Pb2+. This also is a direct analysis because the material being weighed contains the analyte.
Sometimes it is easier to remove the analyte and use a change in mass as the analytical signal. For example, determine a food’s moisture content by a direct analysis. One possibility is to heat a sample of the food to a temperature at which the water in the sample vaporizes. If we capture the vapor in a preweighed absorbent trap, then the change in the absorbent’s mass provides a direct determination of the amount of water in the sample. An easier approach, however, is to weigh the sample of food before and after heating, using the change in its mass as an indication of the amount of water originally present. This technique calls an indirect analysis since we determine the analyte by a signal representing its disappearance.
The indirect determination of moisture content in foods is done by difference. The sample’s initial mass includes the water, whereas the final mass is measured after removing the water. We can also determine an analyte indirectly without its ever being weighed. Again, as with the determination of Pb2+ as PbO2, we take advantage of the analyte’s chemistry. For example, phosphite, PO33–, reduces Hg2+ to Hg22+. In the presence of Cl– a solid precipitate of Hg2Cl2 forms.
2HgCl2 + PO33– + 3H2O « Hg2Cl2 + 2H3O+ + 2Cl– + PO43–
If HgCl2 is added in excess, each mole of PO33– produces one mole of Hg2Cl2. The precipitate’s mass, therefore, provides an indirect measurement of the mass of PO33– present in the original sample.
Procedure
1. The sample is dissolved, if it is not already insoluble.
2. The solution may be treated to adjust the pH (so that the proper precipitate is formed, or to suppress the formation of other precipitates). If it is known that species are present which interfere (by also forming precipitates under the same conditions as the analyte), the sample might require treatment with a different reagent to remove these interferents.
3. The precipitating reagent is added at a concentration that favors the formation of a “good” precipitate (see below). This may require low concentration, extensive heating (often described as “digestion”), or careful control of the pH. Digestion can help reduce the amount of coprecipitation.
4. After the precipitate has formed and been allowed to “digest”, the solution is carefully filtered. The filter is chosen to trap the precipitate; smaller particles are more difficult to filter.
§ Depending on the procedure followed, the filter might be a piece of ashless filter paper in a fluted funnel, or a filter crucible. Filter paper is convenient because it does not typically require cleaning before use; however, filter paper can be chemically attacked by some solutions (such as concentrated acid or base), and may tear during the filtration of large volumes of solution.
§ The alternative is a crucible whose bottom is made of some porous material, such as sintered glass, porcelain or sometimes metal. These are chemically inert and mechanically stable, even at elevated temperatures. However, they must be carefully cleaned to minimize contamination or carryover(cross-contamination). Crucibles are often used with a mat of glass or asbestos fibers to trap small particles.
§ After the solution has been filtered, it should be tested to make sure that the analyte has been completely precipitated. This is easily done by adding a few drops of the precipitating reagent; if a precipitate is observed, the precipitation is incomplete.
5. After filtration, the precipitate – including the filter paper or crucible – is heated. This achieves three purposes:
§ The remaining moisture is removed (drying).
§ Secondly, the precipitate is converted to a more chemically stable form. For instance, calcium ion might be precipitated using oxalate ion, to produce calcium oxalate (CaC2O4); it might then be heated to convert it into the oxide (CaO). It is vital that the empirical formula of the weighed precipitate be known, and that the precipitate be pure; if two forms are present, the results will be inaccurate.
§ The precipitate cannot be weighed with the necessary accuracy in place on the filter paper; nor can the precipitate be completely removed from the filter paper in order to weigh it. The precipitate can be carefully heated in a crucible until the filter paper has burned away; this leaves only the precipitate. (As the name suggests, “ashless” paper is used so that the precipitate is not contaminated with ash.)
6. After the precipitate is allowed to cool (preferably in a desiccator to keep it from absorbing moisture), it is weighed (in the crucible). The mass of the crucible is subtracted from the combined mass, giving the mass of the precipitated analyte. Since the composition of the precipitate is known, it is simple to calculate the mass of analyte in the original sample
Types of Gravimetric Methods
Described above examples illustrate the four gravimetric methods of analysis. If the signal is the mass of a precipitate, the method calls precipitation gravimetry. The indirect determination of PO33– by precipitating Hg2Cl2 is a representative example, as is the direct determination of Cl– by precipitating AgCl.
In electrogravimetry the analyte is deposited as a solid film on one electrode in an electrochemical cell. The oxidation of Pb2+, and its deposition as PbO2 on a Pt anode is one example of electrogravimetry. Reduction also may be used in electrogravimetry. The electrodeposition of Cu on a Pt cathode, for example, provides a direct analysis for Cu2+.
If thermal or chemical energy is used to remove a volatile species, the method calls volatilization gravimetry. In determining the moisture content of food, thermal energy vaporizes the H2O. The amount of carbon in an organic compound may be determined by using the chemical energy of combustion to convert C to CO2.
In particulate gravimetry (isolation method) the analyte is determined following its removal from the sample matrix by filtration or extraction. The determination of suspended solids is one example of particulate gravimetry.
Calculations in Gravimetry
An accurate gravimetric analysis requires that the mass of analyte present in a sample be proportional to the mass or change in mass serving as the analytical signal. For all gravimetric methods this proportionality involves a conservation of mass. For gravimetric methods involving a chemical reaction, the analyte should participate in only one set of reactions, the stoichiometry of which indicates how the precipitate’s mass relates to the analyte’s mass. Thus, for the analysis of Pb2+ and PO33– described earlier, we can write the following conservation equations
Moles Pb2+ = moles PbO2
Moles PO33– = moles Hg2Cl2
Removing the analyte from its matrix by filtration or extraction must be complete. When true, the analyte’s mass can always be found from the analytical signal; thus, for the determination of suspended solids we know that
Filter’s final mass – filter’s initial mass = g suspended solid
whereas for the determination of the moisture content we have
Sample’s initial mass – sample’s final mass = g H2O
Conservation of Mass
Gravimetric calculations are simply an extension of stoichiometric calculations. Our stoichiometric factor is most often based on the amount (in moles) of our analysed species in the material actually weighed.
GF =
GF – gravimetric factor.
FW – formula weight.
http://www.youtube.com/watch?v=dERZhN-01f8
Precipitation Gravimetry
Precipitation gravimetry is based on the formation of an insoluble compound following the addition of a precipitating reagent, or precipitant, to a solution of the analyte. In most methods the precipitate is the product of a simple metathesis reaction between the analyte and precipitant; however, any reaction generating a precipitate can potentially serve as a gravimetric method. Most precipitation gravimetric methods were developed in the nineteenth century as a means for analyzing ores. Many of these methods continue to serve as standard methods of analysis.
A precipitation gravimetric analysis must have several important attributes.
1. The precipitate must be of low solubility, high purity, and of known composition if its mass is to accurately reflect the analyte’s mass. Solubility-product (KSP) must be less than 1×10-8.
2. The precipitate must be in a form that is easy to separate from the reaction mixture and be easy to recover by filtration and washing. From this point of view crystalline precipitates should preferred be used than amorphous.
3. The precipitate must easy and completely transform to weighed form.
4. The precipitate must be unreactive to air, water etc.
5. The precipitate must be something where our analysed species is only a small portion of the precipitate.
Solubility Considerations
An accurate precipitation gravimetric method requires that the precipitate’s solubility be minimal. Many total analysis techniques can routinely be performed with an accuracy of better than ±0.1%. To obtain this level of accuracy, the isolated precipitate must account for at least 99.9% of the analyte. By extending this requirement to 99.99% we ensure that accuracy is not limited by the precipitate’s solubility.
Precipitation is the most important stage of gravimetry. Accuracy of gravimetric methods depends on:
– precipitant choice,
– precipitant amount,
– precipitation conditions.
Calculation of quantity of precipitant
§ In the gravimetric analysis a precipitation is considered practically full if in a solution defined substance is in limits of accuracy of weighing
it is less than 0,0002–
§ Take a volatile precipitant in 2-3 times more than calculated quantities (count on the reaction equation)
§ Take a nonvolatile precipitant in 1,5 times more than calculated quantities (count on the reaction equation)
On precipitation process influence:
1) precipitant amount. Precipitation is complete, when residue amount of analysed species is less than balance sensitivity (
2) temperature. Temperature increasing, frequently, causes precipitate solubility increasing;
3) pH value. Change the pH value represses dissociation of analysed weak electrolytes or dissolves its;
4) complex formation process. Process of complex compounds formation competes with process of precipitate formation. Effect depends on solubility-product value and formation-constant value.
Conditions of precipitation of crystal precipitates:
§ A precipitation are carried from enough diluted solutions by the diluted solution of precipitant.
§ Add a precipitant very slowly, on drops.
§ A solution are mixed continuously by a glass stick to avoid strong local satiation at addition of precipitant.
§ Conduct precipitation from a hot solution, and sometimes heat up also a solution of precipitant (to increase solubility).
§ Filter a precipitate only after cooling of solution.
§ Addition at precipitation of substances which raise solubility of a precipitate (for example, acids).
Conditions of precipitation of amorphous precipitate
§ A precipitation conduct from the concentrated solutions by the concentrated solutions of precipitant.
§ A precipitation conduct from hot solutions.
§ A precipitation conduct in the presence of electrolyte – coagulant.
§ A precipitate quickly filter and do not leave under a matrix solution.
Solubility losses are minimized by carefully controlling the composition of the solution in which the precipitate forms. This, in turn, requires an understanding of the relevant equilibrium reactions affecting the precipitate’s solubility. For example, Ag+ can be determined gravimetrically by adding Cl– as a precipitant, forming a precipitate of AgCl.
Ag+ + Cl– ® AgCl¯ (1)
Figure 1. Solubility of AgCl as a function of pCl. The dashed line shows the predicted SAgCl, assuming that only reaction 1 and equation 2 affect the solubility of AgCl. The solid line is calculated using equation 7, and includes the effect of reactions 3–5. A ladder diagram for the AgCl complexation equilibria is superimposed on the pCl axis.
If this is the only reaction considered, we would falsely conclude that the precipitate’s solubility, SAgCl, is given by
SAgCl = [Ag+] = (2)
and that solubility losses may be minimized by adding a large excess of Cl–. In fact, as shown in Figure 1, adding a large excess of Cl– eventually increases the precipitate’s solubility.
To understand why AgCl shows a more complex solubility relationship than that suggested by equation 2, we must recognize that Ag+ also forms a series of soluble chloro-complexes
Ag+ + Cl– AgCl (3)
Ag+ +2Cl– AgCl2– (4)
Ag+ + 3Cl– AgCl32– (5)
The solubility of AgCl, therefore, is the sum of the equilibrium concentrations for all soluble forms of Ag+.
SAgCl = [Ag+] + [AgCl] + [AgCl2–] + [AgCl32–] (6)
Substituting the equilibrium constant expressions for reactions 3–5 into equation 6 defines the solubility of AgCl in terms of the equilibrium concentration of Cl–.
SAgCl = + K1KSP + b2KSP[Cl–] + b3KSP[Cl–]2 (7)
Equation 7 explains the solubility curve for AgCl shown in Figure 1. As Cl– is added to a solution of Ag+, the solubility of AgCl initially decreases because of reaction 1. Note that under these conditions, the final three terms in equation 7 are small, and that equation 1 is sufficient to describe the solubility of AgCl. Increasing the concentration of chloride, however, leads to an increase in the solubility of AgCl due to the soluble chloro-complexes formed in reactions 3–5.
Also shown in Figure 1 is a ladder diagram for this system. Note that the increase in solubility begins when the higher-order soluble complexes, AgCl2– and AgCl32–, become the dominant species.
Clearly the equilibrium concentration of chloride is an important parameter if the concentration of silver is to be determined gravimetrically by precipitating AgCl. In particular, a large excess of chloride must be avoided.
Mechanism of Precipitation
After the addition of the precipitating agent to the solution of the ion under analysis there is an initial induction period before nucleation occurs. This induction period may range from a very short time period to one which is relatively long, ranging from almost instantaneous to several minutes.
After induction, nucleation occurs, here small aggregates or nuclei of atoms form and it is from these “clumps” of atoms that the crystals which form the filtrate will grow. These nuclei may be composed of just a few atoms each so there may be up to 1010 of the nuclei per mole of precipitating product. As these nuclei form ions from the solution (which at this point are in excess) congregate around them. For example if hydrochloric acid were added very slowly to a solution of silver nitrate, silver chloride nuclei would form and silver ions (which would be in excess relative to Cl- ions) would congregate around them.
In addition to the primary adsorbed silver ion, there are some nitrate ions aggregating further from the AgCl nucleus. These are counter ions and tend to aggregate around the [AgCl:Ag]+ center because these centers have a net positive charge (excess Ag+) and additional negative charge is required to maintain electrical neutrality. The counter ions are less tightly held than the primary adsorbed ions and the counter ion layer is somewhat diffuse and contains ions other than those of the counter ions. These layers of charges are known as the electric double layer.
After nucleation growth occurs, large nuclei grow at the expense of smaller nuclei which dissolve. This process helps produce more easily filtered crystals (since it produces larger crystals).
Growth of larger nuclei or crystallites can be encouraged by digestion, a process which involves heating the solid and mother liquor for a certain period of time. During digestion, small particles dissolve and larger ones grow. Digestion of the product is an important practical process and you will find that most if not all gravimetric analysis involve a digestion period.
http://www.youtube.com/watch?v=dERZhN-01f8
Another important parameter that may affect a precipitate’s solubility is the pH of the solution in which the precipitate forms. For example, hydroxide precipitates, such as Fe(OH)3, are more soluble at lower pH levels at which the concentration of
Ca3(PO4)2 3Ca2+ + 2PO43– (8)
PO43– + H2O HPO42– + OH– (9)
HPO42– + H2O H2PO4– + OH– (10)
H2PO4– + H2O H3PO4 + OH– (11)
Depending on the solution’s pH, the predominate phosphate species is either PO43–, HPO42–, H2PO4–, or H3PO4. The ladder diagram for phosphate, shown in Figure 2a, provides a convenient way to evaluate the pH-dependent solubility of phosphate precipitates. When the pH is greater than 12.4, the predominate phosphate species is PO43–, and the solubility of Ca3(PO4)2 will be at its minimum because only reaction 8 occurs to an appreciable extent (see Figure 2b). As the solution becomes more acidic, the solubility of Ca3(PO4)2 increases due to the contributions of reactions 9–11.
Figure 2. a) Ladder diagram for phosphate; b) Solubility diagram for Ca3(PO4)2 showing the predominate form of phosphate for each segment of the solubility curve.
Solubility can often be decreased by using a nonaqueous solvent. A precipitate’s solubility is generally greater in aqueous solutions because of the ability of water molecules to stabilize ions through solvation. The poorer solvating ability of nonaqueous solvents, even those that are polar, leads to a smaller solubility product. For example, PbSO4 has a Ksp of 1.6 ´10–8 in H2O, whereas in a 50:50 mixture of H2O/ethanol the KSP at 2.6 ´10–12 is four orders of magnitude smaller.
Avoiding Impurities
Precipitation gravimetry is based on a known stoichiometry between the analyte’s mass and the mass of a precipitate. It follows, therefore, that the precipitate must be free from impurities. Since precipitation typically occurs in a solution rich in dissolved solids, the initial precipitate is often impure. Any impurities present in the precipitate’s matrix must be removed before obtaining its weight.
The greatest source of impurities results from chemical and physical interactions occurring at the precipitate’s surface. A precipitate is generally crystalline, even if only on a microscopic scale, with a well-defined lattice structure of cations and anions. Those cations and anions at the surface of the precipitate carry, respectively, a positive or a negative charge as a result of their incomplete coordination spheres. In a precipitate of AgCl, for example, each Ag+ ion in the bulk of the precipitate is bound to six Cl– ions. Silver ions at the surface, however, are bound to no more than five Cl– ions, and carry a partial positive charge (Figure 3).
Figure 3. Schematic model of AgCl showing difference between bulk and surface atoms of silver. Silver and chloride ions are not shown to scale.
Precipitate particles grow in size because of the electrostatic attraction between charged ions on the surface of the precipitate and oppositely charged ions in solution. Ions common to the precipitate are chemically adsorbed, extending the crystal lattice. Other ions may be physically adsorbed and, unless displaced, are incorporated into the crystal lattice as a coprecipitated impurity. Physically adsorbed ions are less strongly attracted to the surface and can be displaced by chemically adsorbed ions.
Sources of coprecipitation
– surface absorption. For example, on surface BaSO4 precipitate adsorbs Pb+ ions;
– occlusion. Coprecipitated impurities placed not on precipitant surface but into precipitate particles. For example, CaC2O4×Na2C2O4;
– mixed crystal formation. Different substances form the same crystal lattices. Ions of impurities built in the units cells of analysed species crystal lattices. For example, KAl(SO4)2×12H2O, KCr(SO4)2×12H2O;
– mechanical entrapment.
One common type of impurity is an inclusion. Potential interfering ions whose size and charge are similar to a lattice ion may substitute into the lattice structure by chemical adsorption, provided that the interferent precipitates with the same crystal structure (Figure 4a). The probability of forming an inclusion is greatest when the interfering ion is present at substantially higher concentrations than the dissolved lattice ion. The presence of an inclusion does not decrease the amount of analyte that precipitates, provided that the precipitant is added in sufficient excess. Thus, the precipitate’s mass is always larger than expected.
Inclusions are difficult to remove since the included material is chemically part of the crystal lattice. The only way to remove included material is through reprecipitation. After isolating the precipitate from the supernatant solution, it is dissolved in a small portion of a suitable solvent at an elevated temperature. The solution is then cooled to re-form the precipitate. Since the concentration ratio of interferent to analyte is lower in the new solution than in the original supernatant solution, the mass percent of included material in the precipitate decreases. This process of reprecipitation is repeated as needed to completely remove the inclusion. Potential solubility losses of the analyte, however, cannot be ignored. Thus, reprecipitation requires a precipitate of low solubility, and a solvent for which there is a significant difference in the precipitate’s solubility as a function of temperature.
Figure 4. Example of coprecipitation:
(a) schematic of a chemically adsorbed inclusion or a physically adsorbed occlusion in a crystal lattice, where C and A represent the cation–anion pair comprising the analyte and the precipitant, and [M] is the impurity;
(b) schematic of an occlusion by entrapment of supernatant solution;
(c) surface adsorption of excess C.
Occlusions, which are a second type of coprecipitated impurity, occur when physically adsorbed interfering ions become trapped within the growing precipitate. Occlusions form in two ways. The most common mechanism occurs when physically adsorbed ions are surrounded by additional precipitate before they can be desorbed or displaced (see Figure 4a). In this case the precipitate’s mass is always greater than expected. Occlusions also form when rapid precipitation traps a pocket of solution within the growing precipitate (Figure 4b). Since the trapped solution contains dissolved solids, the precipitate’s mass normally increases. The mass of the precipitate may be less than expected, however, if the occluded material consists primarily of the analyte in a lower-molecular-weight form from that of the precipitate.
Occlusions are minimized by maintaining the precipitate in equilibrium with its supernatant solution for an extended time. This process is called digestion and may be carried out at room temperature or at an elevated temperature. During digestion, the dynamic nature of the solubility–precipitation equilibrium, in which the precipitate dissolves and re-forms, ensures that occluded material is eventually exposed to the supernatant solution. Since the rate of dissolution and reprecipitation are slow, the chance of forming new occlusions is minimal.
After precipitation is complete the surface continues to attract ions from solution (Figure 4c). These surface adsorbates, which may be chemically or physically adsorbed, constitute a third type of coprecipitated impurity. Surface adsorption is minimized by decreasing the precipitate’s available surface area. One benefit of digestion is that it also increases the average size of precipitate particles. This is not surprising since the probability that a particle will dissolve is inversely proportional to its size. During digestion larger particles of precipitate increase in size at the expense of smaller particles. One consequence of forming fewer particles of larger size is an overall decrease in the precipitate’s surface area. Surface adsorbates also may be removed by washing the precipitate. Potential solubility losses, however, cannot be ignored.
Inclusions, occlusions, and surface adsorbates are called coprecipitates because they represent soluble species that are brought into solid form along with the desired precipitate. Another source of impurities occurs when other species in solution precipitate under the conditions of the analysis. Solution conditions necessary to minimize the solubility of a desired precipitate may lead to the formation of an additional precipitate that interferes in the analysis. For example, the precipitation of nickel dimethylgloxime requires a pH that is slightly basic. Under these conditions, however, any Fe3+ that might be present precipitates as Fe(OH)3. Finally, since most precipitants are not selective toward a single analyte, there is always a risk that the precipitant will react, sequentially, with more than one species.
The formation of these additional precipitates can usually be minimized by carefully controlling solution conditions. Interferents forming precipitates that are less soluble than the analyte may be precipitated and removed by filtration, leaving the analyte behind in solution. Alternatively, either the analyte or the interferent can be masked using a suitable complexing agent, preventing its precipitation.
In some situations the rate at which a precipitate forms can be used to separate an analyte from a potential interferent. For example, due to similarities in their chemistry, a gravimetric analysis for Ca2+ may be adversely affected by the presence of Mg2+. Precipitates of Ca(OH)2, however, form more rapidly than precipitates of Mg(OH)2. If Ca(OH)2 is filtered before Mg(OH)2 begins to precipitate, then a quantitative analysis for Ca2+ is feasible.
Finally, in some cases it is easier to isolate and weigh both the analyte and the interferent. After recording its weight, the mixed precipitate is treated to convert at least one of the two precipitates to a new chemical form. This new mixed precipitate is also isolated and weighed. For example, a mixture containing Ca2+ and Mg2+ can be analyzed for both cations by first isolating a mixed precipitate of CaCO3 and MgCO3. After weighing, the mixed precipitate is heated, converting it to a mixture of CaO and MgO. Thus
Grams of mixed precipitate 1 = grams CaCO3 + grams MgCO3
Grams of mixed precipitate 2 = grams CaO + grams MgO
Although these equations contain four unknowns (grams CaCO3, grams MgCO3, grams CaO, and grams MgO), the stoichiometric relationships between CaCO3 and CaO
Moles CaCO3 = moles CaO
and between MgCO3 and MgO
Moles MgCO3 = moles MgO
provide enough additional information to determine the amounts of both Ca2+ and Mg2+ in the sample.
Controlling Particle Size
Following precipitation and digestion, the precipitate must be separated from the supernatant solution and freed of any remaining impurities, including residual solvent. These tasks are accomplished by filtering, rinsing, and drying the precipitate. The size of the precipitate’s particles determines the ease and success of filtration. Smaller, colloidal particles are difficult to filter because they may readily pass through the pores of the filtering device. Large, crystalline particles, however, are easily filtered.
By carefully controlling the precipitation reaction we can significantly increase a precipitate’s average particle size. Precipitation consists of two distinct events: nucleation, or the initial formation of smaller stable particles of precipitate, and the subsequent growth of these particles. Larger particles form when the rate of particle growth exceeds the rate of nucleation.
A solute’s relative supersaturation, RSS, can be expressed as
RSS = (12)
where Q is the solute’s actual concentration, S is the solute’s expected concentration at equilibrium, and Q – S is a measure of the solute’s supersaturation when precipitation begins. A large, positive value of RSS indicates that a solution is highly supersaturated. Such solutions are unstable and show high rates of nucleation, producing a precipitate consisting of numerous small particles. When RSS is small, precipitation is more likely to occur by particle growth than by nucleation.
Examining equation 12 shows that we can minimize RSS by either decreasing the solute’s concentration or increasing the precipitate’s solubility. A precipitate’s solubility usually increases at higher temperatures, and adjusting pH may affect a precipitate’s solubility if it contains an acidic or basic anion. Temperature and pH, therefore, are useful ways to increase the value of S. Conducting the precipitation in a dilute solution of analyte, or adding the precipitant slowly and with vigorous stirring are ways to decrease the value of Q.
There are, however, practical limitations to minimizing RSS. Precipitates that are extremely insoluble, such as Fe(OH)3 and PbS, have such small solubilities that a large RSS cannot be avoided. Such solutes inevitably form small particles. In addition, conditions that yield a small RSS may lead to a relatively stable supersaturated solution that requires a long time to fully precipitate. For example, almost a month is required to form a visible precipitate of BaSO4 under conditions in which the initial RSS is 5.
An increase in the time required to form a visible precipitate under conditions of low RSS is a consequence of both a slow rate of nucleation and a steady decrease in RSS as the precipitate forms. One solution to the latter problem is to chemically generate the precipitant in solution as the product of a slow chemical reaction. This maintains the RSS at an effectively constant level. The precipitate initially forms under conditions of low RSS, leading to the nucleation of a limited number of particles. As additional precipitant is created, nucleation is eventually superseded by particle growth. This process is called homogeneous precipitation.
Two general methods are used for homogeneous precipitation. If the precipitate’s solubility is pH-dependent, then the analyte and precipitant can be mixed under conditions in which precipitation does not occur. The pH is then raised or lowered as needed by chemically generating
CO(NH2)2 + H2O « CO2 + 2NH3
NH3 + H2O « NH4+ +
The hydrolysis of urea is strongly temperature-dependent, with the rate being negligible at room temperature. The rate of hydrolysis, and thus the rate of precipitate formation, can be controlled by adjusting the solution’s temperature. Precipitates of BaCrO4, for example, have been produced in this manner.
In the second method of homogeneous precipitation, the precipitant itself is generated by a chemical reaction. For example, Ba2+ can be homogeneously precipitated as BaSO4 by hydrolyzing sulphamic acid to produce SO42–.
NH2SO3H + 2H2O « NH4+ + H3O+ + SO42–
Homogeneous precipitation affords the dual advantages of producing large particles of precipitate that are relatively free from impurities. These advantages, however, may be offset by increasing the time needed to produce the precipitate, and a tendency for the precipitate to deposit as a thin film on the container’s walls. The latter problem is particularly severe for hydroxide precipitates generated using urea.
Figure 5. Schematic model of the solid–solution interface at a particle of AgCl in a solution containing excess AgNO3.
An additional method for increasing particle size deserves mention. When a precipitate’s particles are electrically neutral, they tend to coagulate into larger particles. Surface adsorption of excess lattice ions, however, provides the precipitate’s particles with a net positive or negative surface charge. Electrostatic repulsion between the particles prevents them from coagulating into larger particles.
Consider, for instance, the precipitation of AgCl from a solution of AgNO3, using NaCl as a precipitant. Early in the precipitation, when NaCl is the limiting reagent, excess Ag+ ions chemically adsorb to the AgCl particles, forming a positively charged primary adsorption layer (Figure 5). Anions in solution, in this case NO3– and OH–, are attracted toward the surface, forming a negatively charged secondary adsorption layer that balances the surface’s positive charge. The solution outside the secondary adsorption layer remains electrically neutral. Coagulation cannot occur if the secondary adsorption layer is too thick because the individual particles of AgCl are unable to approach one another closely enough.
Coagulation can be induced in two ways: by increasing the concentration of the ions responsible for the secondary adsorption layer or by heating the solution. One way to induce coagulation is to add an inert electrolyte, which increases the concentration of ions in the secondary adsorption layer. With more ions available, the thickness of the secondary absorption layer decreases. Particles of precipitate may now approach one another more closely, allowing the precipitate to coagulate. The amount of electrolyte needed to cause spontaneous coagulation is called the critical coagulation concentration.
Heating the solution and precipitate provides a second way to induce coagulation. As the temperature increases, the number of ions in the primary adsorption layer decreases, lowering the precipitate’s surface charge. In addition, increasing the particle’s kinetic energy may be sufficient to overcome the electrostatic repulsion preventing coagulation at lower temperatures.
Filtering the Precipitate
After precipitation and digestion are complete, the precipitate is separated from solution by filtration using either filter paper or a filtering crucible.
filter paper
filtering crucible
Preparing your Filter Paper
Folding a piece of filter paper for insertion into a conical filter consists of a simple set of steps shown here in the six photographs below. From left to right and top to bottom, one first folds the round piece of filter paper in half and creases it. Then it is folded again and creased to produce a quarter circle. One outer layer of paper is separated from the other three (not two and two!) and the opening made wider by squeezing slightly together at the creases. The conical shaped piece of filter paper is placed into a glass or plastic funnel and wetted slightly with distilled water from your wash bottle:
The most common filtering medium is cellulose-based filter paper, which is classified according to its filtering speed, its size, and its ash content on ignition. Filtering speed is a function of the paper’s pore size, which determines the particle sizes retained by the filter. Filter paper is rated as fast (retains particles > 20–25 mm), medium fast (retains particles >
Filter paper is hygroscopic and is not easily dried to a constant weight. As a result, in a quantitative procedure the filter paper must be removed before weighing the precipitate. This is accomplished by carefully igniting the filter paper. Following ignition, a residue of noncombustible inorganic ash remains that contributes a positive determinate error to the precipitate’s final mass. For quantitative analytical procedures a low-ash filter paper must be used. This grade of filter paper is pretreated by washing with a mixture of HCl and HF to remove inorganic materials. Filter paper classed as quantitative has an ash content of less than 0.010% w/w. Qualitative filter paper typically has a maximum ash content of 0.06% w/w.
Procedure for filtering through a filtering crucible. The trap is used to prevent water from a water aspirator from backwashing into the suction flask.
An alternative method for filtering the precipitate is a filtering crucible. The most common is a fritted glass crucible containing a porous glass disk filter. Fritted glass crucibles are classified by their porosity: coarse (retaining particles > 40–60 mm), medium (retaining particles > 10–15 mm), and fine (retaining particles > 4–5.5 mm). Another type of filtering crucible is the Gooch crucible, a porcelain crucible with a perforated bottom. A glass fiber mat is placed in the crucible to retain the precipitate, which is transferred to the crucible in the same manner described for filter paper. The supernatant is drawn through the crucible with the assistance of suction from a vacuum aspirator or pump.
Rinsing the Precipitate
Filtering removes most of the supernatant solution. Residual traces of the supernatant, however, must be removed to avoid a source of determinate error. Rinsing the precipitate to remove this residual material must be done carefully to avoid significant losses of the precipitate. Of greatest concern is the potential for solubility losses. Usually the rinsing medium is selected to ensure that solubility losses are negligible. In many cases this simply involves the use of cold solvents or rinse solutions containing organic solvents such as ethanol. Precipitates containing acidic or basic ions may experience solubility losses if the rinse solution’s pH is not appropriately adjusted. When coagulation plays an important role in determining particle size, a volatile inert electrolyte is often added to the rinse water to prevent the precipitate from reverting into smaller particles that may not be retained by the filtering device. This process of reverting to smaller particles is called peptization. The volatile electrolyte is removed when drying the precipitate.
When rinsing a precipitate there is a trade-off between introducing positive determinate errors due to ionic impurities from the precipitating solution and introducing negative determinate errors from solubility losses. In general, solubility losses are minimized by using several small portions of the rinse solution instead of a single large volume. Testing the used rinse solution for the presence of impurities is another way to ensure that the precipitate is not overrinsed. This can be done by testing for the presence of a targeted solution ion and rinsing until the ion is no longer detected in a freshly collected sample of the rinse solution. For example, when Cl– is known to be a residual impurity, its presence can be tested for by adding a small amount of AgNO3 to the collected rinse solution. A white precipitate of AgCl indicates that Cl– is present and additional rinsing is necessary. Additional rinsing is not needed, however, if adding AgNO3 does not produce a precipitate.
Choice of a rinsings liquid
§
§ Amorphous precipitates are rinsed by solutions of volatile electrolytes to avoid of peptization of a precipitate
§ Precipitates with high solubility are rinsed by solutions of electrolytes which contain the same ion with a precipitate
Concentration of impurities Сn which remained in a precipitate afterrinsings:
Drying the Precipitate
§ Drying in porcelain and glass filtering crucibles
!!! Drying of weighed form are leaded to its constant weight, that is the difference between its parallel weighing will not exceed ±0,0002 г
drying box
muffle furnace
Finally, after separating the precipitate from its supernatant solution the precipitate is dried to remove any residual traces of rinse solution and any volatile impurities. The temperature and method of drying depend on the method of filtration, and the precipitate’s desired chemical form. A temperature of
Filter paper’s ability to absorb moisture makes its removal necessary before weighing the precipitate. This is accomplished by folding the filter paper over the precipitate and transferring both the filter paper and the precipitate to a porcelain or platinum crucible. Gentle heating is used to first dry and then to char the filter paper. Once the paper begins to char, the temperature is slowly increased. Although the 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 charred the temperature is slowly raised to a higher temperature. At this stage any carbon left 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
Gravimetric Sulfate Determination
The small heating pads available in the laboratory are large enough to hold the three 400 mL beakers to be used in this experiment, as shown in the photo at the right. Heating all three in such close quarters has its dangers, of course, mainly that of tipping one over and losing one sample! So as not to mix up your samples, label each 1,2 and 3 with a graphite pencil.
Notice below how one of the beakers containing the dissolved sulfate sample is placed beneath a buret filled with 5% barium chloride solution. The precipitate is formed by adding the barium chloride slowly to the hot sulfate solution. The slow addition assures that the relative supersaturation always remains low so as to avoid the formation of a non-crystalline colloidal suspension of barium sulfate particles. The aluminum foil seen in this photograph is not necessary as the temperature of the solution is not high enough to melt the paint on the base of the ring stand.
The addition of barium chloride solution can also be done dropwise using an eyedropper (below right) but you must make sure that the eyedropper is clean before you start, that you do the addition using a measured amount of barium chloride solution placed in a small beaker and that you don’t allow any barium chloride solution to contaminate the rubber bulb of the eyedropper. Using this method allows you to add the barium chloride solution to the beakers while they are continuing to be heated on the hot pads.
The filtration will be carried out using glass or plastic funnels fitted with ashless filter paper. Ashless filter paper is pure cellulose and will decompose in the presence of heat and air to water and carbon dioxide. No residual non-volatile substances remain. Ashless filter paper comes in the form of circles which must be folded appropriately to trap all of the barium sulfate precipitate.
The funnel rack (below left) is useful to allow you to carry out the filtration of all samples simultaneously. Make sure that the filter paper you use is ASHLESS. Erlenmeyer flasks may be used to catch the filtrate. As each flask fills up, examine the filtrate to make sure that it is free of any precipitate. If the filtrate samples are clear they may be discarded. Presence of precipitate suggests either a hole in your filter paper or overfilling a funnel so that precipitate was not caught by the filter. In either case, ask your instructor about the procedure to be followed.
As is often the case in laboratory chemistry there is more than one correct way to carry out a procedure and the same can be said for the incorrect execution of a procedure. Here are four pictures to illustrate the point. The first two illustrate correct ways of pouring your filtrate and suspended precipitate into the filter in the funnel. In the first one on the right note that the student is using a stirring rod to provide a route for the solution into the filter. The lip of the beaker is sufficiently close to the filter so that the falling solution never speeds up fast enough to form droplets and end up above the filter paper or, worse, on the lab bench.
The picture at the left shows the pouring operatioear the end of the transfer. A stirring rod is used again but held in place across the top of the beaker for stability. The lip of the beaker is again close enough to the funnel not to let any of the solution get away. The only criticism that might be offered here is that the length of the fall of solution is a little too great. So as to minimize the possibility of loss owing to solution turbulence as it falls, the path of fall ought to be shortened.
The next picture on the right shows a student simply pouring his solution out of the beaker. Immediately after this photo was taken, the falling liquid curled up onto the beaker owing to the surface tension of water. Some was lost on the filter above the filter paper and some dribbled down the side of the beaker. A stirring rod ought to be used to provide a good path for the solution being transferred.
The photo immediately to the left shows the use of a stirring rod to provide a path, but the rod is at an angle of inclination too low to provide a good path for the solution. The path ought to be more vertical or there is a good chance that some solution will be lost by dribbling off the rod. In addition, the distance of the lip of the beaker is too far from the funnel and there is every likelihood that at some point some of the liquid traveling down the rod will separate from it and be lost from the analysis.
There certainly is no need to beat this subject to death, but I’ll lay two more photos on you then move on to the rest of these helpful hints. Even though it is slightly out of focus, the photo on the left shows a nearly vertical stirring rod. This is about as good as it is going to get. The possibility that any of the supernatant is going to dribble down the outside of the beaker under the lip is kept to a minimum by this technique. Please be cognizant that the rod DOES NOT TOUCH the filter paper. You must avoid that so as not to make holes through which the precipitate might be lost. The photo on the right illustrates that one ought to try to pour practically all of the supernatant through the filter paper before starting on the precipitate. That way the filter paper remains unclogged by precipitate until the very end. The student is in the process of transferring the precipitate. This can be done by a combination of low-volume washes from your wash bottle and coaxing the precipitate out with your rubber policement, as shown here.
Finally, the setup on the right shows a filtrate which resulted from leakage either through or around the filter paper in the small beaker. The appropriate technique would have been for the student to fold a new piece of filter paper and place it UNDER the first one and continue filtering, including a refiltration of the filtrate with the residue followed by washing as described below and then ashing of both pieces of filter paper. In fact the student washed the precipitate off the filter paper back into the suspension yet to be filtered and threw away the old filter paper. Recovery of the filter paper from the trash can is not recommended because of the possibility it has picked up other material, so he has the choice of risking a low result if all precipitate was not washed from it or starting over. In any case, the filtrate in the larger beaker is appropriately clear, so there is no apparent leakage of precipitate into it.
Hot water washes of 15 mL each need to be tested with two drops of silver nitrate solution for any lingering chloride ion. These three photographs, from left to right show what you might expect to see as all of the chloride is leached from the precipitate.
Fold your filter paper so that the precipitate is trapped inside and so that the final shape is small enough to be stuffed into a clean crucible which has been previously weighed. It doesn’t matter if the filter paper is still damp. It may be easier to fold and to stuff if it IS damp, but be careful not to fold it in a manner which causes the paper itself to rupture because you might lose some of your precipitate. Note in the image on the left that the filter paper has been stuffed into the crucible and heating has begun. The image in the middle shows that as the ashing process progresses the filter paper blackens and there is the possibility of spontaneous combustion. Keep your crucible cover ready to extinguish any flame which might produce sufficient turbulence to allow the escape of fragments of filter paper coated with precipitate. Finally, the image on the right shows the crucible just prior to final ashing.
By way of some final observations, as the filter paper is heated and loses its moisture and then begins to distill destructively, turning to carbon and releasing the inevitable tars from the thermal destruction of the cellulose, it is important to cover it so as not to allow it to burst into flame because the turbulence might allow some barium sulfate caked on it to escape with the fumes.
On the other hand, covering the crucible has its perils as well as the photo on the left shows. Look closely and you can see that the crucible is held in mid-air because the top has stuck to the rim. The destructively distilled tar can cause the cap to stick to the crucible allowing both to be lifted together by the cover knob. Don’t lift such a pair too high because the cover may suddenly separate from the crucible causing breakage of the crucible.
Allow the crucible and cover to cool, then gently pry the two apart. The rim of the crucible and the roof of the cover both will have tar and carbonaceous material caked on them. Note the underside of the crucible cap in the photo at the left
Both can be subjected to heating and ultimately lose all of that residue as shown in the photo immediately on the right. If there is any suspicion that some of the barium sulfate may have deposited on the roof of the cover, you may weigh it after heating it thoroughly to remove all of the carbonaceous material , then wash it, heat it to dryness and weigh it again. Experience has shown though that rarely does any barium sulfate deposit with the destructively distilled residue on the cover.
Even if you have a speck of black material, it is good to continue heating for a while to get rid of it if it is carbonaceous material as shown here in the photos at the left and the right.
The carbonaceous material on the underside of a cover can be quickly removed by heating it upside down. The one at the left was heated as shown and the carbonaceous material oxidized and disappeared in a matter of minutes.
Composition of Final Precipitate
Composition of Final Precipitate
The quantitative application of precipitation gravimetry, 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 standard gravimetric method for the determination of magnesium involves the precipitation of MgNH4PO4 ´ 6H2O. Unfortunately, this precipitate is difficult to dry at lower temperatures without losing an inconsistent amount of hydrated water and ammonia. Instead, the precipitate is dried at temperatures above
An additional problem is encountered when the isolated solid is nonstoichiometric. For example, precipitating Mn2+ as Mn(OH)2, followed by heating to produce the oxide, frequently produces a solid with a stoichiometry of MnOx, where x varies between 1 and
Weight of the precipitate’s form for different types of precipitates
|
Precipitate type |
Molar weight , g/mol |
Weight of precipitate, g |
|
Amorphous (Fe2O3 × H2O) |
до »100 |
0,07 – 0,10 |
|
Easy crystal (CaCO3) |
до » 100 |
0,10 – 0,15 |
|
Heavy crystal (BaCO3) |
» 200 |
0,2 – 0,4 |
|
Very heavy crystal (PbSO4) |
» 300 |
0,5 |
Weight of gravimetric form:
§ For crystal precipitates –
§ For amorphous precipitates –
!!!! It is necessary to remember: the more weight defined substance, the above relative accuracy of results of the analysis.
Evaluating Precipitation Gravimetry
Scale of Operation. The scale of operation for precipitation gravimetry is governed by the sensitivity of the balance and the availability of sample. To achieve an accuracy of ±0.1% using an analytical balance with a sensitivity of ±0.1 mg, the precipitate must weigh at least 100 mg. As a consequence, precipitation gravimetry is usually limited to major or minor analytes, and macro or meso samples. The analysis of trace level analytes or micro samples usually requires a microanalytical balance.
Accuracy. For macro–major samples, relative errors of 0.1–0.2% are routinely achieved. The principal limitations are solubility losses, impurities in the precipitate, and the loss of precipitate during handling. When it is difficult to obtain a precipitate free from impurities, an empirical relationship between the precipitate’s mass and the mass of the analyte can be determined by an appropriate standardization.
Precision. The relative precision of precipitation gravimetry depends on the amount of sample and precipitate involved. For smaller amounts of sample or precipitate, relative precisions of 1–2 ppt are routinely obtained. When working with larger amounts of sample or precipitate, the relative precision can be extended to several parts per million. Few quantitative techniques can achieve this level of precision.
Sensitivity. For any precipitation gravimetric method, we can write the following general equation relating the signal (grams of precipitate) to the absolute amount of analyte in the sample
Grams precipitate = k ´ grams of analyte (13)
where k, the method’s sensitivity, is determined by the stoichiometry between the precipitate and the analyte. Note that equation 13 assumes that a blank has been used to correct the signal for the reagent’s contribution to the precipitate’s mass.
For example, the determination of Fe as Fe2O3. Using a conservation of mass for Fe we write
2 ´ moles Fe2O3 = moles Fe
Converting moles to grams and rearranging yields an equation in the form of 13
g Fe2O3 = ´ g Fe
where k is equal to
k = (14)
As can be seen from equation
Selectivity. Due to the chemical nature of the precipitation process, precipitants are usually not selective for a single analyte. For example, silver is not a selective precipitant for chloride because it also forms precipitates with bromide and iodide. Consequently, interferents are often a serious problem that must be considered if accurate results are to be obtained.
Volatilization Gravimetry
A second approach to gravimetry is to thermally or chemically decompose a solid sample. The volatile products of the decomposition reaction may be trapped and weighed to provide quantitative information. Alternatively, the residue remaining when decomposition is complete may be weighed. In thermogravimetry, which is one form of volatilization gravimetry, the sample’s mass is continuously monitored while the applied temperature is slowly increased.
Whether the analysis is direct or indirect, volatilization gravimetry requires that the products of the decomposition reaction be known. This requirement is rarely a problem for organic compounds for which volatilization is usually accomplished by combustion and the products are gases such as CO2, H2O, and N2. For inorganic compounds, however, the identity of the volatilization products may depend on the temperature at which the decomposition is conducted.
Thermogravimetry. The products of a thermal decomposition can be deduced by monitoring the sample’s mass as a function of applied temperature. (Figure 9). The loss of a volatile gas on thermal decomposition is indicated by a step in the thermogram. The change in mass at each step in a thermogram can be used to identify both the volatilized species and the solid residue.
Figure 9. Thermogram for CaC2O4 ´H2O.
Once the products of thermal decomposition have been determined, an analytical procedure can be developed. For example, the thermogram in Figure 9 shows that a precipitate of CaC2O4´H2O must be heated at temperatures above
Equipment. Depending on the method, the equipment for volatilization gravimetry may be simple or complex. In the simplest experimental design, the weight of a solid residue is determined following either thermal decomposition at a fixed temperature or combustion. Thermal decomposition or combustion is accomplished using a Bunsen or Meker burner, a laboratory oven or a muffle furnace, with the volatile products vented to the atmosphere. The weight of the sample and the solid residue are determined using an analytical balance.
Constant-temperature decomposition or combustion followed by trapping and weighing the volatilized gases, requires more specialized equipment. Decomposition of the sample is conducted in a closed container, and the volatilized gases are carried by a purge-gas stream through one or more selective absorbent traps.
In a thermogravimetric analysis, the sample is placed in a small weighing boat attached to one arm of a specially designed electromagnetic balance and placed inside an electric furnace. The temperature of the electric furnace is slowly increased at a fixed rate of a few degrees per minute, and the sample’s weight is monitored.
Qualitative Applications of Gravimetry
Although not in common use, precipitation gravimetry still provides a reliable means for assessing the accuracy of other methods of analysis or for verifying the composition of standard reference materials.
Inorganic Analysis. The most important precipitants for inorganic cations are chromate, the halides, hydroxide, oxalate, sulphate, sulphide, and phosphate. A summary of selected methods, grouped by precipitant, is shown in Table 1. Many inorganic anions can be determined using the same reactions by reversing the analyte and precipitant. For example, chromate can be determined by adding BaCl2 and precipitating BaCrO4. Methods for other selected inorganic anions are summarized in Table 2. Methods for the homogeneous generation of precipitants are shown in Table 3.
The majority of inorganic precipitants show poor selectivity. Most organic precipitants, however, are selective for one or two inorganic ions. Several common organic precipitants are listed in Table 4.
Table 1. Selected Gravimetric Method for Inorganic Cations Based on Precipitation
Table 2. Selected Gravimetric Methods for Inorganic Anions Based on Precipitation
Table 3. Reactions for the Homogeneous Preparation of Selected Inorganic Precipitants
Advantage of organic precipitants consists in the following:
§ Solubility of precipitate with organic precipitants is less.
§ Precipitates with organic reagents are crystal.
§ Precipitates with organic reagents are purer as on their surface impurity are less adsorbed.
§ Organic precipitant have higher selectivity and specificity.
§ The gravimetric factor at definition with organic reagents on much less so, accuracy of definition increases.
Table 4. Selected Gravimetric Methods for Inorganic Cations
Based on Precipitation with Organic Precipitants
Requirements to precipitants:
§ It is desirable, that a precipitant was volatile compound.
§ A precipitant should be specific – to precipitate a defined ion in the presence of others ions.
Organic Analysis. Several organic functional groups or heteroatoms can be determined using gravimetric precipitation methods; examples are outlined in Table 5. Note that the procedures for the alkoxy and alkimide functional groups are examples of indirect analyses.
Quantitative Applications of Gravimetry
Unlike precipitation gravimetry, which is rarely used as a standard method of analysis, gravimetric methods based on volatilization reactions continue to play an important role in chemical analysis.
Inorganic Analysis. Determining the inorganic ash content of organic materials, such as polymers and paper, is an example of a direct volatilization gravimetric analysis. The sample is weighed, placed in an appropriate crucible, and the organic material is carefully removed by combustion. The crucible containing the residue is then heated to a constant weight using either a burner or an oven.
Another example of volatilization gravimetry is the determination of dissolved solids in water and wastewater. In this method a sample of the water is transferred to a weighed dish and dried to a constant weight at either 103–105 °C, or at
Table 5. Selected Gravimetric Methods for the Analysis of Organic Functional Groups and Heteroatoms Based on Precipitation
Indirect analyses based on the weight of the residue remaining after volatilization are commonly used in determining moisture in a variety of products and in determining silica in water, wastewater, and rocks. Moisture is determined by drying a preweighed sample with an infrared lamp or in a low-temperature oven. The difference between the original weight and the weight after drying equals the mass of water lost.
The determination of silicon is commonly encountered in metallurgical and mining laboratories responsible for the analysis of ores, slags, and alloys. The volatilization gravimetric method, which is appropriate for samples containing high concentrations of silicon. The determination of carbon in steels and other metal alloys can be determined by heating the sample. The carbon is converted to CO2, which is collected in an appropriate absorbent trap, providing a direct measure of the amount of C in the original sample.
Organic Analysis. The most important application of volatilization gravimetry to the analysis of organic materials is an elemental analysis. When burned in a stream of pure O2, many elements, such as carbon and hydrogen, are released as gaseous combustion products, such as CO2 and H2O. The combustion products are passed through preweighed tubes containing appropriate absorbents. The increase in the mass of these tubes provides a direct indication of the mass percent of carbon and hydrogen in the organic material.
Alkaline metals and earths in organic materials can be determined by adding H2SO4 to the sample before combustion. Following combustion, the metal remains behind as a solid residue of metal sulfate. Silver, gold, and platinum can be determined by burning the organic sample, leaving a metallic residue of Ag, Au, or Pt. Other metals are determined by adding HNO3 before combustion, leaving a residue of the metal oxide.
Volatilization gravimetry is also used to determine biomass in water and wastewater. Biomass is a water quality index, providing an indication of the total mass of organisms contained within a sample of water. A known volume of the sample is passed through a preweighed 0.45-mm membrane filter or a glass-fiber filter and dried at
Use gravimetry in the pharmaceutical analysis
– Determination of moisture in drugs (indirect volatilization gravimetry).
– Determination of water (direct volatilization gravimetry).
– Determination of the dry rest in extracts, tinctures (indirect volatilization gravimetry).
– Determination of sulphatic ashes and ashes (particulate gravimetry).
– Determination of drugs (precipitation and particulate gravimetry).
Definition of the dry residue in tinctures.
In weighed weighing bottle or flat-bottomed dish (preliminary finished to constant weight) introduce 2,0 ml of the tincture to be examined. Evaporate to dryness on a water-bath and dry in an oven at 100-
Definition loss on drying of drugs.
In weighed weighing bottle or flat-bottomed dish (preliminary finished to constant weight) introduce 1,0 ml of the substance to be examined (lactose, glucose, analginum). Dry in an oven at 100-
Definition of the general ashes.
Use Method I unless otherwise directed in the monograph.
Method I
For vegetable drugs
Incinerate 2 to
For other substances
Carry out the above method using
Method II
Heat a silica or platinum crucible to redness for 30 min, allow to cool in a desiccator and weigh. Unless otherwise prescribed, evenly distribute
Definition of acid-insoluble ash.
Use Method I unless otherwise directed in the monograph.
Method I
Boil the ash for 5 minutes with 25 ml of 2M hydrochloric acid, collect the insoluble matter in a sintered-glass crucible or on an ashless filter paper, wash with hot water and ignite. Calculate the percentage of acid-insoluble ash with reference to the air-dried drug.
Method II
Ash insoluble in hydrochloric acid is the residue obtained after extracting the sulphated or total ash with hydrochloric acid, calculated with reference to
To the crucible containing the residue from the determination of sulphated or total ash, add 15 ml of water R and 10 ml of hydrochloric acid R, cover with a watch-glass, boil the mixture gently for 10 min and allow to cool. Filter through an ashless filter, wash the residue with hot water R until the filtrate is neutral, dry, ignite to dull redness, allow to cool in a desiccator and weigh. Reheat until the difference between 2 consecutive weighings is not more than 1 mg.
Definition of sulphatic ash.
Use Method I unless otherwise directed.
Method I
Heat a platinum dish to redness for 10 minutes, allow to cool in a desiccator and weigh. Unless otherwise specified in the monograph, place
Method II
Ignite a suitable crucible (for example, silica, platinum, porcelain or quartz) at 600 ±
If the amount of the residue so obtained exceeds the prescribed limit, repeat the moistening with sulphuric acid R and ignition, as previously, for 30 min periods until 2 consecutive weighings do not differ by more than 0.5 mg or until the percentage of residue complies with the prescribed limit.
Definition of Iron
Iron can be determined gravimetrically by precipitating as Fe(OH)3 and igniting to Fe2O3. The sample to be analyzed is weighed and transferred to a 400-mL beaker where it is dissolved in 50 mL of H2O and 10 mL of
Fe2+ + NO3– + 2H+ = Fe3+ + NO2 + H2O
After boiling to remove the oxides of nitrogen, the solution is diluted to 200 mL, brought to boiling, and Fe(OH)3 is precipitated by slowly adding 1:1 NH3 until the odor of NH3 is detected.
Fe3 + + 3NH3 + 3H2O FFe (OH) 3¯ + 3NH4 +.
The solution is boiled for an additional minute, and the precipitate is allowed to settle to the bottom of the beaker. The precipitate is then filtered and washed with several portions of hot 1% w/v NH4NO3 until no Cl– is found in the wash water. Finally, the precipitate is ignited to constant weight at 500–550 °C, and weighed as Fe2O3.
2Fe (OH) 3 Fe2O3 + 3H2O.
