Why digest the precipitate

Many coagulated precipitates do not respond well to washing with distilled water because as the second electric layer is removed excess Cl - for example the first remains on all particles with an electric charge of the same sign. The result is that there is a return to the repulsive state and an effective increase in the radius of the particles which then begin once again to separate as colloidal particles.

The process is called peptization and is to be avoided if some of the precipitate is not to be lost. One way around this for many precipitates is to encourage digestion by heating and also by increasing the electrolyte concentration by washing with a reagent which will go off as a gas during the drying process. In choosing such a wash, it is imperative that the procedure has been carried out and has been shown to yield reproducible, quantitative results.

Unexpected side reactions, complex formation and changes in solubility with added reagents are sufficiently unpredictable to make intuition in the absence of experience unacceptable. During the precipitation procedure a number of other problems can arise to give erroneous positive or negative results. Among these are surface adsorption, mixed crystal formation, occlusion and mechanical entrapment. Any ions may be carried down during a precipitation as the result of surface adsorption.

Likewise, the ions Cl - and SO 4 2 - can have the same effect. In the quantitative determination of some transition metals, iron for example as Fe OH 3 , zinc, cadmium and manganese may be present as impurities and all three form sparingly soluble hydroxides as well, though each with greater solubility than the hydroxide of iron:.

Mixed crystal formation can occur if two ions have the same charge, if their ionic diameters are sufficiently close to fit into the same crystal lattice. Ions which commonly interfere with each other are shown in the table below with their ionic diameters in picometers given after each.

In cases where one has a known interference of one ion with the other it is necessary to find methods of removing one before carrying out a precipitation of the other, or using a precipitating reagent in which there is no interference.

Occlusion and mechanical entrapment. If a precipitation procedure is carried out too quickly, pockets of solvent and spectator ions can form, trapping them within the precipitate particles and dashing one's hope of removing them during the washing procedure.

This is another reason why the relative supersaturation must be kept as low as possible so that in principle at least, all precipitation occurs only at the surface of a growing solid particle, devoid of solvent pockets. All of these problems of coprecipitation of unwanted ions can lead to positive or negative errors.

In the case of mixed crystal formation, the direction of the error depends on the relative atomic weight of the ion which replaces that which is desired in the precipitate. In the case of the precipitation of zinc hydroxide, mixed crystal formation with manganese would produce a negative error but with cadmium or zinc a positive error. A solution containing a reagent which produces a desired ion to effect precipitation, often by gentle heating of the solution, offers an exquisite means for obtaining well-formed large crystal particles which lend themselves splendidly to the technique of filtration.

The model we use to explain why this happens also uses the concept of relative supersaturation. The initial nucleation of sparingly soluble particles offers a surface template which favors "locking" onto ions in the vicinity which by the luck of the draw and the kinetic molecular theory find themselves at the right energy and orientation to enter the crystal lattice.

Ions isolated from a growing crystal are not favored to enter this process because at least two are required, both at the right energy and orientation to start the growth of a new crystal. If the concentration of one ion of a sparingly soluble salt increases gradually by slow homogeneous synthesis in a solution, then as its concentration reaches the threshold of supersaturation for the ion pair, a relatively small number of nucleated particles grows to larger size because the probability of finding a place in an existing crystal lattice for any single ion is greater than that of a spontaneous creation a new crystal from dissolved and randomly arranged ions rather than a large number of nucleated particles growing in constant competition with the rest and thus remaining small.

The result for the latter is a non-filterable precipitate, but one in the former which filters quite well. See the demonstration of this effect at The resulting precipitate must be heated until a stable dry state is reached. Some understanding of typical precipitate properties is mandatory for repeatable results to be achieved. Note in the figure at the right that whereas AgCl achieves a stable dry weight just above o C, BaSO 4 does not do so until it reaches a temperature in the vicinity of o C Aluminum oxide, Al 2 O 3 , loses water slowly as the temperature rises to o C at which point it achieves stability.

Some compounds decompose in several stages, reaching stable plateaus. Between o C and o C it slowly decomposes to CaO where it remains stable until its melting point at o C. A device not seen often in analytical laboratories but useful for producing automatic plots of mass of sample vs. The amount of NH 3 is crucial to this procedure. This increases the mass of the ignited precipitate, and gives a positive determinate error.

Each mole of Mg 2 P 2 O 7 contains two moles of magnesium and each mole of Mg PO 3 2 contains only one mole of magnesium. One mole of Mg 2 P 2 O 7 weighs Two moles of Mg PO 3 2 weigh Two additional steps in the procedure help to form a precipitate that is free of impurities: digestion and reprecipitation.

This is done for the same reason that the precipitation is carried out in an ammonical solution; using dilute ammonia minimizes solubility losses when we rinse the precipitate.

Although no longer a common analytical technique, precipitation gravimetry still provides a reliable approach for assessing the accuracy of other methods of analysis, or for verifying the composition of standard reference materials.

In this section we review the general application of precipitation gravimetry to the analysis of inorganic and organic compounds. The majority of inorganic precipitants show poor selectivity for the analyte. Many organic precipitants, however, are selective for one or two inorganic ions.

The precipitate is digested at 80—90 o C for at least two hours. Ashless filter paper pulp is added to the precipitate to aid in its filtration. After filtering, the precipitate is ignited to constant weight at o C. Alternatively, the precipitate is filtered through a fine porosity fritted glass crucible without adding filter paper pulp , and dried to constant weight at o C. This procedure is subject to a variety of errors, including occlusions of Ba NO 3 2 , BaCl 2 , and alkali sulfates.

Other standard methods for the determination of sulfate in water and wastewater include ion chromatography see Chapter 12 , capillary ion electrophoresis see Chapter 12 , turbidimetry see Chapter 10 , and flow injection analysis see Chapter Several organic functional groups or heteroatoms can be determined using precipitation gravimetric methods. Note that the determination of alkoxy functional groups is an indirect analysis in which the functional group reacts with and excess of HI and the unreacted I — determined by precipitating as AgCl.

The stoichiometry of a precipitation reaction provides a mathematical relationship between the analyte and the precipitate. Because a precipitation gravimetric method may involve additional chemical reactions to bring the analyte into a different chemical form, knowing the stoichiometry of the precipitation reaction is not always sufficient. Even if you do not have a complete set of balanced chemical reactions, you can use a conservation of mass to deduce the mathematical relationship between the analyte and the precipitate.

The following example demonstrates this approach for the direct analysis of a single analyte. To determine the amount of magnetite, Fe 3 O 4 , in an impure ore, a 1. Filtering, rinsing, and igniting the precipitate provides 0. A conservation of mass requires that the precipitate of Fe 2 O 3 contain all iron originally in the sample of ore. In one analysis, the zinc in a The copper in a separate A conservation of mass requires that all zinc in the alloy is found in the final product, Zn 2 P 2 O 7.

We know there are 2 moles of Zn per mole of Zn 2 P 2 O 7 ; thus. Because we can precipitate each analyte selectively, finding their respective concentrations is a straightforward stoichiometric calculation.

But what if we cannot separately precipitate the two analytes? To find the concentrations of both analytes, we still need to generate two precipitates, at least one of which must contain both analytes. Although this complicates the calculations, we can still use a conservation of mass to solve the problem. Igniting the precipitate converts it to a mixture of Al 2 O 3 and MgO that weighs 1. The masses of the solids provide us with the following two equations.

With two equations and four unknowns, we need two additional equations to solve the problem. Substituting the equations for g MgO and g Al 2 O 3 into the equation for the combined weights of MgO and Al 2 O 3 leaves us with two equations and two unknowns. For aluminum, we find that. A sample of a silicate rock that weighs 0. The mixture of chloride salts is dissolved in a mixture of ethanol and water, and treated with HClO 4 , precipitating 0.

A conservation of mass requires that all the potassium originally in the KCl ends up in the KClO 4 ; thus. Given the mass of KClO 4 , we use the third equation to solve for the mass of KCl in the mixture of chloride salts. The previous problems are examples of direct methods of analysis because the precipitate contains the analyte. In an indirect analysis the precipitate forms as a result of a reaction with the analyte, but the analyte is not part of the precipitate.

As shown by the following example, despite the additional complexity, we still can use conservation principles to organize our calculations. An impure sample of Na 3 PO 3 that weighs 0. After digesting, filtering, and rinsing the precipitate, 0. This is an example of an indirect analysis because the precipitate, Hg 2 Cl 2 , does not contain the analyte, Na 3 PO 3. Although you can write the balanced reactions for any analysis, applying conservation principles can save you a significant amount of time!

A redox reaction must obey a conservation of electrons because all the electrons released by the reducing agent, Na 3 PO 3 , must be accepted by the oxidizing agent, HgCl 2. Knowing this, we write the following stoichiometric conversion factors:. Now we are ready to solve the problem. As you become comfortable using conservation principles, you will see ways to further simplify problems.

For example, a conservation of electrons requires that the electrons released by Na 3 PO 3 end up in the product, Hg 2 Cl 2 , yielding the following stoichiometric conversion factor:. This conversion factor provides a direct link between the mass of Hg 2 Cl 2 and the mass of Na 3 PO 3. After we isolate the precipitate by filtration, we dissolve it in acid and precipitate and weigh the molybdate as PbMoO 3.

Suppose we know that our sample is at least What is the minimum amount of sample that we need for each analysis? Finally, we convert this mass of Na 3 PO 4 to the corresponding mass of sample. A sample of 0. If a sample contains more than A precipitation reaction is a useful method for identifying inorganic and organic analytes.

Because a qualitative analysis does not require quantitative measurements, the analytical signal is simply the observation that a precipitate forms. Although qualitative applications of precipitation gravimetry have been replaced by spectroscopic methods of analysis, they continue to find application in spot testing for the presence of specific analytes [Jungreis, E. Spot Test Analysis ; 2nd Ed. The scale of operation for precipitation gravimetry is limited by the sensitivity of the balance and the availability of sample.

As a consequence, precipitation gravimetry usually is limited to major or minor analytes, in macro or meso samples. The analysis of a trace level analyte or a micro sample requires a microanalytical balance. For a macro sample that contains a major analyte, a relative error of 0. The principal limitations are solubility losses, impurities in the precipitate, and the loss of precipitate during handling. For a smaller amount of sample or precipitate, a relative precision of 1—2 ppt is obtained routinely.

When working with larger amounts of sample or precipitate, the relative precision extends to several ppm. Few quantitative techniques can achieve this level of precision. For any precipitation gravimetric method we can write the following general equation to relate the signal grams of precipitate to the absolute amount of analyte in the sample.

Consider, for example, the determination of Fe as Fe 2 O 3. In other words, it helps to form a precipitate with the largest possible formula weight. We can also improve sensitivity by forming a precipitate that contains fewer units of the analyte.

Suppose you wish to determine the amount of iron in a sample. Of the three choices, the greatest sensitivity is obtained with Fe 2 O 3 because it provides the largest value for k. Due to the chemical nature of the precipitation process, precipitants usually are not selective for a single analyte.

For example, silver is not a selective precipitant for chloride because it also forms precipitates with bromide and with iodide. Interferents often are a serious problem and must be considered if accurate results are to be obtained. Equipment needs are few—beakers, filtering devices, ovens or burners, and balances—inexpensive, routinely available in most laboratories, and easy to maintain. Theory and Practice All precipitation gravimetric analyses share two important attributes.

Avoiding Impurities In addition to having a low solubility, a precipitate must be free from impurities. Ball-and-stick model showing the lattice structure of AgCl. A silver ion on the surface, for example, carries a partial positive charge. Three examples of impurities that may form during precipitation.

The cubic frame represents the precipitate and the blue marks are impurities present as a inclusions, b occlusions, and c surface adsorbates. Inclusions are randomly distributed throughout the precipitate. To identify the cation, the color of the precipitate and its solubility in excess are noted.

Similar processes are often used to separate chemically similar elements, such as the rare earth metals. Digestion, or precipitate ageing , happens when a freshly-formed precipitate is left, usually at a higher temperature, in the solution from which it is precipitated.

It results in cleaner and bigger particles. Coprecipitation is the carrying down by a precipitate of substances normally soluble under the conditions employed.

It is an important issue in chemical analysis, where it is often undesirable, but in some cases it can be exploited. In gravimetric analysis , it is a problem because undesired impurities often coprecipitate with the analyte, resulting in excess mass.

On the other hand, in the analysis of trace elements, as is often the case in radiochemistry , coprecipitation is often the only way of separating an element. Category : Chemical processes. Read what you need to know about our industry portal chemeurope.

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