Digestion

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STEPS OF A GRAVIMETRY ANALYSIS

Preparation of the Sample
Digestion
Precipitation
Filtration and Washing
Calculation
Weighing
Drying or Ignition
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• 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
• To maintain low solubility of the precipitate.
• To obtain the precipitate in a form suitable for
  filtration.
• Proper adjustment of the solution condition may
  also mask potential interferences.

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• FACTORS THAT MUST BE CONSIDERED WHEN PREPARING
  THE SOLUTION
• Volume of the solution during precipitation.
• Concentration range of the test substance.
• The presence and concentrations of other
  constituents.
• Temperature
• pH
• Formation and Properties of Precipitates
• Analyte Precipitating Agent ?
  Supersaturation ? Precipitation
• Two steps are involved in precipitation
• Nucleation and Particle Growth
• The particle size of a precipitate is determined
  by the faster of these two steps.
• In nucleation a few atoms, ions or molecules join
  together to give a stable solid called nuclei.

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Further precipitation then involves a competition
between additional nucleation and particle growth
(deposition of ions/molecules on the surface of
existing nuclei). If Rate of Nucleation gt Rate
of Particle Growth Precipitate containing a large
number of small particles. Rate of Nucleation lt
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.
So, which condition we want Rate of Nucleation gt
Rate of Particle Growth OR Rate of Nucleation lt
Rate of Particle Growth
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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).
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• 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.

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• 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.

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

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• 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)2CO 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.

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• IMPURITIES IN PRECIPITATES
• Occurred in two ways
• ? Co-Precipitation
• A process in which normally soluble compounds are
  carried out of solution by a precipitate
• ? 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
• Mixedcrystal formation
• Occlusion and mechanical entrapment

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• 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 opposed 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.

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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.
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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.
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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.
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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
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• Solid solution formed by interferent (Pb2)
  substituting for lattice ion (Ba2).
• 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.

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• OCCLUSION AND MECHANICAL ENTRAPMENT
• 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.
• Methods to minimize/avoid this thing happen are
  to remove the interferences prior to
  precipitation or to select a different reagent.

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• Two methods to prevent occlusions are
• 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.
• Digestion and aging the precipitate.

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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.
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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.
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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.
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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.
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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.
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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.
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• 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.

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The GF, is the weight of analyte per unit weight
of precipitate.
a, b mole of analyte and precipitate.
Weight A weight precipitate x GF
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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/analyte appearing in the numerator and
denominator must be chemically equivalent.
Sought, A (analyte) Weight ,S (precipitate) Granimetric Factor GF
Br AgBr
I AgI
Fe Fe2O3
Fe3O4 Fe2O3
P Mg2P2O7
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Sought, A (analyte) Weight ,S (precipitate) Gravimetric Factor GF
P2O5 Mg2P2O7
Na5P3O10 Mg2P2O7
S BaSO4
SO2 BaSO4
Ba BaSO4
Ni Ni(C4H7N2O2)2
Al Al2O3
NaB4O7.10H2O B2O3
CaCO3 CaC2O4
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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 the percentage of P2O5 in the
sample 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)
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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)
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Exercise 0.3516 g phosphate detergent was burnt
to lose the organic compounds. It was then added
into a hot HCl to convert P (FW 30.974) to
H3PO4.6H2O. After filtration, the precipitate was
then converted to Mg2P2O7 (FW 222.57) by
ignition. The weight of the remaining compound is
0.2161 g. Calculate P in the sample.