Filtration and Washing STEPS IN GRAVIMETRIC ANALYSIS Preparation of the Sample Precipitation Digestion Filtration and Washing Calculation STEPS of a GRAVIMETRY ANALYSIS Weighing Drying or Ignition

Preparation of the solution   Solid sample must be dissolved in a suitable solvent. Some form of preliminary separation may be necessary to eliminate interfering materials. The Purposes of Solution Preparation i) To maintain low solubility of the precipitate. ii) To obtain the precipitate in a form suitable for filtration. iii) Proper adjustment of the solution condition may also mask potential interferences.

Factors that Must be Considered when Preparing the Solution   i) Volume of the solution during precipitation. ii) Concentration range of the test substance. iii) The presence and concentrations of other constituents. iv) Temperature v) pH

Formation and Properties of Precipitates   Analyte + Precipitating Agent  Supersaturation Precipitation Two steps are involved in precipitation: 1. Nucleation 2.Particle Growth The particle size of a precipitate is determined by which one is faster between these two steps. In nucleation a few atoms, ions or molecules join together to give a stable solid called nuclei (nucleus). Further precipitation then involves a competition between additional nucleation and particle growth (the process where more ions are added to the nucleus to form colloids with sizes in the range of 1-100nm in diameter).  

If: In high supersaturated solution…. Rate of Nucleation > Rate of Particle Growth Precipitate containing a large number of small particles. Colloidal is formed in the solution. In less supersaturated solution…… Rate of Nucleation < Rate of Particle Growth Precipitate containing a smaller number of larger particles is produced.   Precipitates with large particle are more easily handled during filtration and washing. ** In general, a dilute solution with low supersaturation enhances particles growth that results in large particle size.

Factors That Determine the Particle Size of precipitates   Von Weimarn introduced the concept of relative supersaturation, The particle size of precipitates is inversely proportional to the relative supersaturation of the solution during precipitation. Relative Supersaturation = Q - S S Where, Q = concentration of the mixed reagent before precipitation S = solubility of the precipitate at equilibrium (QS) is a measure of the degree of supersaturation. The rate of nucleation and the rate of particle growth depend on the supersaturation (QS).

If: (Q-S) Becomes too Great High relative supersaturation, nucleation is favored, many small particles (Colloidal precipitates form).   (Q-S) is Small The smaller will be the relative supersaturation, particle growth will predominate, few but larger particles size (Crystalline precipitates form). To minimize supersaturation and obtain large crystals, conditions should be adjusted so that Q will be as low as possible (Q ) and S will be relatively large (S ) during precipitation.

Several steps are commonly taken   Precipitation from dilute solution, (Q ). Adding precipitating reagents slowly with effective stirring. Stirring prevents local excesses of the reagent, (Q  ) Precipitation from hot solution. The solubility of precipitates increases with temperature, (S). After the precipitate was formed, the particle size can be improved by digestion process or by precipitation from homogeneous solution.

The effect of relative supersaturation on the particle sizes of precipitate is shown in the diagram below.

Precipitation from Homogeneous Solution   A technique in which a precipitating agent is generated in a solution of the analyte by slow chemical reaction. The precipitating ion is not added to the solution but is slowly generated throughout the solution by a homogeneous chemical reaction. The advantages  No locally excesses of precipitating agent.  The supersaturation (Q-S) is kept low at all the time, so that the precipitate obtained is much more dense and free from impurities than precipitates formed by the conventional method.  Substances that ordinarily precipitate only as amorphous solid frequently precipitate from homogenous solution as well-formed crystals.

(NH2)2 CO + 3H2O  CO2 + 2NH4+ + 2OH Example:  Urea (NH2CONH2)   Used for generation of hydroxide ion. Hydrolysis of urea will increase the pH (more alkaline). The ammonia slowly liberated raises the pH of the solution, and react with metal ions to form metal hydroxides (or hydrous oxides) precipitate. (NH2)2 CO + 3H2O  CO2 + 2NH4+ + 2OH Process is controlled by heating the solution just below 100oC and a 1-2 hr heating period is needed to complete the typical precipitation. Used for the precipitation of Al, Ga, Th, Bi, Fe and Sn as hydroxides.

Impurities In Precipitates   Occurred in two ways:  Co-Precipitation Known as absorbed and adsorbed impurities. Co-precipitation is a process where the impurity is precipitated along with the desired precipitate, even though the solubility of the impurity has not been exceeded.  Post Precipitation Foreign compound precipitates on top of the desired precipitate. For example, post precipitation of magnesium oxalate occurs if a precipitate of calcium oxalate is allowed to stand too long before being filtered.

Co-precipitation  3 types of coprecipitation: Surface adsorption Mixed–crystal formation Occlusion and mechanical entrapment Surface Adsorption   A process in which a normally soluble compound is carried out of solution on the surface of a precipitate. Adsorption is a reversible process because it is accompanied by the opposite process of desorption. The two opposing processes lead to a state of dynamic equilibrium known as adsorption equilibrium.

The position of the adsorption equilibrium depends on numerous factors;   Effects of Surface Area The amount of a substance adsorbed is directly proportional to the total surface area of the adsorbent. Effect of Concentration Adsorption of ions increases with an increase of their concentration (not proportional). Effect of Temperature Adsorption is an exothermic process. Rise of temperature means decreased adsorption. Effects of the Nature of the Adsorbed Ions Adsorbents with ionic crystal lattices prefer to adsorb ions, which form sparingly soluble or common ions with the precipitate. For example, BaSO4 precipitate prefers to adsorb its own common ions, Ba2+ or SO42- dependent on which is present in the solution in excess.

Occurs with all precipitates especially which has a very large surface area like colloidal particle.   The surfaces of the precipitate contain some primary adsorbed ion, either the lattice cation or the lattice anion (the excess lattice ion). If AgCl is precipitated by the addition AgNO3 to excess NaCl, Cl will be adsorbed on the precipitate surface. The lattice ion adsorbed is called the primary layer or primary adsorbed ion. The surface of the precipitate has a minus charge (because of the Cl). The particles carry the same charge on the surface and because of that they repel each other and will not easily coagulate to form larger particles. To balance this charge, the adsorbed ions will also attract an ion of opposite charge (called the counter ion) such as Na+, which then surround the precipitate particles. For a silver halide precipitate, AgX, two cases are possible.

A colloidal silver chloride particle suspended in; (a) Silver nitrate solution (excess Ag+) (b) KCl (excess Cl-)

If nitrate is co-precipitated, the results will be too high because nitrate weights more than chloride.   A lower weight counter ion will result in the weight of the precipitate being too low.

Methods for minimizing adsorbed Impurities    Washing Thorough washing of a precipitate removes surface impurities or replaces them with adsorbed substances that are volatile on drying.  Digestion A process in which a precipitate is heated (without stirring) for an hour or more in the solution from which it was formed (the mother liquor).  Reprecipitation A drastic but effective way to minimize the effects of adsorption. This approach is one of last resort since it significantly add to the analysis time.

Inclusion   A type of co-precipitation in which interferences are incorporated into the precipitating crystal either; (a) As substitute elements on a lattice site or, (b) As extra elements between sites called interstitial. When two salts or compounds have the same type of formula crystallizing in similar geometric forms (called isomorphous), one compound can replace part of the other in a crystal. The result is the formation of mixed crystals. ·                   Example: Co precipitation of PbSO4 in BaSO4 (b)  Ba2+ SO42 Ba2+ SO42 Ba2+ Ba2+ SO42 Ba2+ SO42 K+ SO42 Ba2+ SO42– Pb2+ SO42 SO42 Ba2+ SO42 Ba2+ Pb2+ SO42 Ba2+ SO42 Ba2+ Ba2+ SO42 Ba2+ SO42

(a) Solid solution formed by interferent (Pb2+) substituting for lattice ion (Ba2+).   (b) Solid solution formed by interferent (K+) filling an interstitial site in lattice. Inclusions occur throughout the crystal, not only on the surface; therefore, changes in particle size will not affect the extent of inclusion.

Occlusion and Mechanical Entrapment   Methods to minimize/avoid are to remove the interferences prior to precipitation or to select a different reagent. If crystal growth is too rapid, some counter ions do not have time to escape from the surface so it becomes trapped or occluded within a growing crystal. In another type of occlusion; mechanical entrapment, two crystals grow together and trap a species in the space between them. These crystals lie close together during growth. Two methods to prevent occlusions are: 1) Slow precipitation by slow addition of dilute reagent to a dilute solution of the analyte, gives counter ions time to leave and helps break up pockets. 2) Digestion and aging the precipitate.

Occlusion and Mechanical Entrapment

Digestion of the Precipitate   A process in which a precipitate is heated (without stirring) for an hour or more in the solution from which it was formed (the mother liquor) . It will produce larger and purer particles. The small particles tend to dissolve (less supersaturation compared to the solution before) and reprecipitate on the surface of larger particles/crystals causing the particles to grow even larger. During this process, water is expelled from the solid to give a denser mass that has a smaller specific surface area for adsorption. Digestion is usually done at elevated temperature to speed the process. Heating tends to decreases the number of adsorbed particles and the effective charge in the adsorbed layer, thereby aiding coagulation.

Filtering And Washing the Precipitates   A precipitate may be separated by filtering it through: Ashless Filter Paper – leave very little ash when it burnt. Sintered Glass – able to stand temperature up to 500oC. Sintered Porcelain – able to stand temperature up to 1000oC. The choice of a medium depends on the type of precipitate and on the temperature at which the precipitate is to be heated. Curdy precipitates are usually filtered with sintered-glass filtering crucibles. Crystalline precipitates are filtered through filter paper. The purpose of washing a precipitate is to remove adsorbed impurities and the mother liquor.

For crystalline precipitates, several rinsings with small amounts of pure water will remove the impurities. Washing coagulated colloids with pure water will dilute and remove foreign ions and the counter ions will occupy a larger volume. The charge of the primary ions repel each other and revert back to the colloidal form. This process is called peptization and results in the loss of colloidal precipitate through the filter. Coagulated colloid are commonly washed with electrolyte solution (e. g. HNO3, NH4OH) for AgCl but not KNO3 because it is nonvolatile.

To test the completeness of washing:   Example: Cl is determined by precipitating with AgNO3 reagent. The filtrate is tested for Ag+ by adding NaCl or dilute HCl. If the solution becomes cloudy it means the washing process is not complete yet. Washing is continued until the test is negative.

Drying (Heating) and Ignition of Precipitates   Filtered precipitates normally are heated or ignited to produce a constant weight of precipitate. Drying removes the solvent(s) and any volatile species carried down with the precipitate. The drying can be done by heating at 110oC to 120oC for 1 to 2 hours (for the precipitate which is in a form suitable for weighing). If a precipitate must be converted to a more suitable form for weighing, ignition at a much higher temperature is required.

Effect of Temperature on Calcium Oxalate, CaC2O4 Mass

Around 200oC, calcium oxalate exists as the hydrate CaC2O4.H2O. ii) Around 400oC, calcium oxalate exists as CaC2O4. CaC2O4.H2O  CaC2O4 + H2O iii) Around 700oC, calcium oxalate is converted to CaCO3. CaC2O4  CaCO3 + CO iv) Above 1000oC, calcium carbonate is converted to CaO. CaCO3  CaO + CO2   Each compound has different drying/ignition behavior which must be determined on a case by case basis.

Cooling and Weighing Precipitates   After heated, the precipitates must be cooled to room temperature in a desiccator before weighing. The process of heating, cooling and weighing is repeated until a constant weight is achieved or the difference between two consecutive weighing is not more than 0.0002 g.

Calculating the Results   The result of a gravimetric determination usually is reported as a percentage of analyte Two values are needed: the weight of the analyte and the weight of the sample. To calculate the weight of analyte from the weight of the precipitate a gravimetric factor (GF) is used.

Examples and Calculation In Gravimetry The GF, is the weight of analyte per unit weight of precipitate.   a, b: mole of analyte and precipitate.

Two points should be noted in setting up a gravimetric factor:   The molecular weight (or atomic weight) of the analyte is in the numerator; that of the substance weighed is in the denominator. The number of molecules or atoms appearing in the numerator and denominator must be chemically equivalent. Example: The phosphorus in a sample of phosphate rock weighing 0.5428 g is precipitated as MgNH4PO4.6H2O and ignited to Mg2P2O7. If the ignited precipitate weighs 0.2234g, calculate: (a)         the percentage of P2O5 in the sample (b)         the weight of the precipitate of MgNH4PO4.6H2O. (FW P2O5 = 141.94, FW Mg2P2O7 = 222.55, FW MgNH4PO4.6H2O = 245.40) (Ans. (a)26.25, (b) 0.4927)

Exercise:   Magnetite is an ore with a formula Fe3O4 or FeO.Fe2O3. 1.1423 g magnetite ore was solubilized in a concentrated HCl. The Fe2+ and Fe3+ solutions are all converted to Fe3+ using HNO3. Fe3+ was precipitated as Fe2O3.xH2O with the addition of NH3. After washing, the filtrate was ignited at a very high temperature producing 0.5394 g pure Fe2O3. Calculate the %Fe and % Fe3O4 in the sample. (33.32%; 46.04%) 0.3516 g phosphate detergent was burnt to lose the organic compounds. It was then added into a hot HCl to convert P (30.974) to H3PO4.6H2O. After filtration, the precipitate was then converted to Mg2P2O7 (222.57) by ignition. The weight of the remaining compound is 0.2161 g. Calculate %P in the sample.

Methods for Prevention and Elimination of the Three Co-precipitation