Gravimetric analysis

Gravimetric analysis
Gravimetry is a general term applied to methods of analysis in
which the determination are based upon measurements of
Gravimetric analysis is the process of isolating and weighing of
the final product with known, pure, stable and definite
chemical structure.
The separation of the element or the compound containing it,
may be done in a number of ways, the most important of
which are: i)
Separation by volatilization
Separation by extraction or partition
Separation by adsorption and ion exchange
Separation by precipitation.
Our interest will be devoted to the separation by precipitation
Gravimetry by precipitation:
Precipitation methods are based on precipitating the
substance to be determined from solution as an insoluble
compound of known chemical composition. The content of the
component is then calculated from the weight of the resulting
precipitate. The precipitate is often ignited before weighing, to
decompose and convert it into a substance of a different
composition, which is more suitable to be weighed.
These processes are generally involved in any gravimetric
1. Precipitation.
2. Filtration and washing
3. Drying or/and ignition.
The precipitation process takes place in three stages. In the
first one, ions collide with each other in the supersaturated
solution to form primary nuclei. Superimposition of ions on the
primary nuclei to form colloidal particles is the second stage,
and in the third stage colloidal particles aggregate to form
visible precipitates which according to the concentration of
the solutions used, rate of nuclei formation and the nature of
the precipitated material may be crystalline (as barium
sulphate), curdy (as silver chloride) or gelatinous (as ferric
Stage 1
ions in supersaturated solution
Primary nuclei
(non filterable)
Stage 2
Colloidal particals
Fine crystals
Stage 3
Final form
Coarse crystals
Crystalline aggregate
e.g. BaSO4
Stabilized Colloid
Colloidal aggregate
Gel (e.g. Fe(OH)3)
Curd (e.g. AgCl)
Choice of precipitant
One of the important considerations in gravimetric analysis is
the choice of the right precipitating agent. Specific reagents
would react with one single species of compounds thus
preventing the precipitation of the other contaminant neither
by true nor by co-precipitation. However the ideal of specific
reagents for one substance has not been realized. Reagents
are mostly selective, i.e. react with a group of compounds.
However selectivity can be improved by adjusting the
condition of the experiment: In some cases change of the
oxidation number of some members of the selectivity. For
example reduction of ferric ion to ferrous prevents its
precipitation as hydroxide with ammonia. The masking effect of
certain complexing agents may be also useful.
Fundamental Requirements for the precipitation process.
Three fundamental requirements must be fulfilled during
precipitation process.
1. The substance to be determined must be precipitated
quantitatively. In practice, this usually means that the
quantity remaining in solution does not exceed the
minimum detectable by the ordinary analytical
balance, viz 0.1 mg
2. The precipitate must be pure not contaminated (this
point will be discussed in detail later)
3. The precipitate must be in a physical form suitable for
Precipitates of relatively large crystals are very convenient
1. They are retained readily by the filter paper,
2. Their surface is not extensive, accordingly, they do not
readily adsorb impurities from the solution, and
3. They are easily washed.
Precipitates consisting of very small crystals, such as BaSO4 or
CaC2O4 are less convenient in this respect.
Amorphous precipitates, especially if gelatinous, such as
Al(OH)3, have extensive specific surface and therefore adsorb
considerable amounts of impurities which are difficult to wash
Requirements for the weighed form
1. Its composition should correspond exactly to its
chemical formula. For example, the Fe(OH)3 formed in
gravimetric determination of iron does not correspond
exactly to the formula Fe(OH)3, but contains variable
amounts of water and which are not known exactly.
Therefore, it will be more correct to write its formula as
(Fe2O3 n.H2O). When ferric hydroxide is ignited, all this
water is removed and a compound of quite definite
composition is formed, exactly corresponding the
formula Fe2O3. Moreover, when a precipitate is ignited
the water and any volatile impurities retained by it, are
completely removed.
2. The weighed form must have adequate chemical
stability. For example, CaO precipitate readily absorbs
H2O and CO2 from air; therefore it is sometimes
converted into CaSO4 by treatment with sulphuric acid
and excess acid is removed by evaporation.
3. Finally, the content of the element being determined
in the precipitate should be as low as possible.
The colloidal state:
The colloidal state of matter is distinguished by a certain range
of particle size (0.001 to 0.1) as a consequence of which
certain characteristic properties become apparent. Ordinary
filter papers will retain particles up to a diameter of 10, so the
colloidal solutions in this respect behave like true solutions.
Properties of colloids:
1. The solid is dispersed in the form of aggregates of the ions
or molecules while in true solutions the dissolved solid is
homogeneously dispersed in the form of individual ions or
2. In contrast to the true solution, the solid present in a
colloidal suspension has a negligible effect on such
properties as freezing point, boiling point, and osmotic
3. The colloidal suspension have the tendency to scatter
visible radiation. This when a light beam passes through a
colloidal solution, its path can be readily seen (Tyndall
4. Owing to the smallness of the size of the particles, the
ratio of surface to weight is extremely large. Phenomena
which depend upon the size of the surface, such as
adsorption, will play therefore an important part with
substances in the colloidal state.
5. Typical colloidal solutions may remain without forming a
precipitate for a very long time. This clearly proves the
existence of certain factors preventing cohesion of the
colloidal particles. One of these factors, is the presence
of like electric charges on the colloidal particles; this
cause mutual repulsion of the particles, which are
prevented from joining into large aggregates.
We shall consider a case in which a colloidal arsenious
sulphide is formed by the use of H2S in acid medium as a
precipitating agent. Here sulphide ions are primarily adsorbed
(since every precipitate has a tendency to adsorb its own
ions), and some hydrogen ions are secondarily adsorbed have
been termed counter ions. Thus the so called electrical double
layer is set up between the particles and the solution.
If a small amount of electrolytes is added to a colloidal
aggregates that sink to the bottom of the vessel. The
coagulation effect is explained simply with the assumption that
the colloidal particles adsorb oppositely charged ions from
electrolyte to these primarily adsorbed. Accordingly, the
charges on the colloidal particles are decreased so much that
they can join together. Thus coagulation effects of electrolytes
increase rapidly with increasing the valency of the ions of the
opposite charge to that on the colloidal particles.
Upon washing the precipitate with water, part of the adsorbed
electrolyte is removed. The electrolyte concentration in the
supernatant liquid may fall below the coagulation value
flocculation), the precipitate may pass into colloidal solution
The formation of amorphous precipitates:
In this case addition of each portion of precipitant causes
rapid formation of an enormous number of very minute
crystalline nuclei in the liquid; these grow not by deposition of
the substance on their surfaces but by joining to form larger
aggregates which sink by gravity to the bottom of the vessel.
Amorphous precipitates have:
1. Enormous surface area and therefore adsorb readily
various extraneous substances from solutions.
2. Since the bonds between the individual crystal nuclei in
the aggregates are relatively weak, the aggregates
may break up again to give a colloidal solution.
Substances of very low solubility form amorphous precipitate
for example metal sulphides and hydroxides. Because of the
very low solubility, the solubility product is greatly exceeded by
addition of even small amounts of precipitant, and this favors
the rapid formation of numerous crystal nuclei.
The type of
precipitate formed depends not only on the individual
conditions. For example BaSO4 is precipitated in crystalline form
out of dilute aqueous solution. However, if it is precipitated out
of mixture of water with 30 - 60% of alcohol, which greatly
reduces the solubility of BaSO4, a colloidal solution or an
amorphous precipitate is formed.
Homogeneous precipitation
When a precipitant is added to a solution, even when the
solution is dilute and well stirred, there will always be some local
regions of high supersaturation, and impure precipitate forms.
However, if the precipitant is added directly, but is rather
generated slowly by homogeneous chemical reaction within
the solution at a rate comparable to the rate of particle
growth, the extent of supersaturation will not reach as high a
value as could exist if the precipitant was added directly. This
technique is called homogeneous precipitation, and its
principal requirements are that the chemical used must not
cause any reaction during its addition to the system, any that it
can be made to from the precipitant by a slow reaction, often
a hydrolysis reaction.
That homogeneous formation of crystalline precipitates will
lead to less occlusion than the classical methods seems
reasonable when it is recalled that occlusion is believed to
result from envelopment of adsorbed ions and counter ions by
the rapidly growing crystalline solid. But during homogeneous
precipitation crystal growth occur slowly and there is, therefore,
more time for the adsorbed ions to be replaced on the surface
by the lattice ions. But contamination by isomorphic and
nonisomorphic inclusion is not appreciably by homogenous
The best known example of homogeneous precipitation is the
use of the hydrolysis of urea to generate hydroxyl ions for the
precipitation of aluminium, ferric iron, and other heavy metal
ions. The precipitant (hydroxyl ion) is formed by the following
reaction which takes place by heating urea just below the
boiling point of water until the pH of the solution rises to the
desired value:
(NH2)2 CO + 3 H2O  CO2 + 2NH4+ + 2 OHThe homogeneous precipitated heavy metal hydroxide is
superior to that precipitated by the classical method; in that it is
much purer, and more dense, compact and easily filterable.
Many applications of the above technique involves the
homogeneous removal of hydrogen ion in order to increase the
concentration of an anion acting as the precipitant for acid
soluble compounds such as calcium oxalate. Here, the calcium
and oxalate ions are brought together in a solution that is
sufficiently acidic to prevent formation of the solid. Precipitation
is then induced through homogeneous neutralization of the
acid by the ammonium hydroxide liberated by the hydrolysis of
urea. Large crystals result, free from contamination with
magnesium or phosphate after one precipitation.
Homogeneous precipitation of sulphates, oxalates, phosphates
and sulphides can be made by adding the corresponding alkyl
ester and boiling to hydrolyse the ester into the inorganic anion,
i.e., the precipitant. In this way, calcium and magnesium can
be precipitated quantitatively as the oxalate, barium and
strontium as the sulphate and aluminium and zinc as the
phosphate. The generation reactions are represented by the
following equation:
+ 2 H2 O 
2 C2H5OH
+ H2C2O2
+ 2 H2 O 
2 CH3OH + 2 H+ + SO42-
+ 3 H2 O 
3 C2H5OH + H3PO4
CH3 CS NH2 +
H2 O
Cation-release method
Each of the proceeding examples involves the slow generation
of an anion to unite with the sought-for cation forming a
precipitate. Some methods of homogeneous precipitation
entail the reverse procedure that is the slow generation of a
cation for the precipitation of the desired anion. One well
known example is the liberation of barium ion, in the presence
of sulphate ion (the anion to be determined by precipitation),
through the reaction between hydrogen peroxide and
ethylenediamine tetra-acetic acid (EDTA). The barium-EDTA
complex is soluble and is very stable, but the organic ligand is
destructively oxidized by hydrogen peroxide, and barium ion,
freed from the complex, is then able to react with the sulphate
2. Filtration
For separating the precipitate from the mother liquor, filtration
is the most common operation. Essentially, filtration is effected
by passing the slurry, through a porous medium. Its techniques
involve the use of filter-paper or filter crucibles and the choice
of one technique or the other is decided by the type of
precipitate and the subsequent treatment it may need.
A) Filtration on paper:
Filter papers of various sizes and porosities are commercially
available. The porosity of the paper used is chosen according
to the nature of the precipitate. Three grades are available in
the fine precipitates as of barium sulphate. No 41 papers have
wider pores and are suitable for gelatinous precipitates as
hydrated ferric and aluminium oxides. The more open texture of
whatman No. 40 papers makes them suitable for well
crystallized precipitates as that of magnesium ammonium
hydrochloric and hydrofluoric acids to remove silicon and other
minerals which cannot be volatilized during the subsequent
ignition, so as to minimize the ash left after the paper has been
burnt away. Such acid-treated papers are called ashless filterpapers; they usually leave, after incineration, not more than 1
gm. residue per paper.
Paper pulp prepared by shaking ashless paper clippings with
hot water can be used as a filtering medium for the rapid
filtration of the most finely divided precipitates.
B) Filtering crucibles:
When the precipitate as to be weighed after drying only, or
then it cannot be ignited in presence of paper, a filtering
crucible is useful.
It also permits the use of suction to increase the speed filtration
determinations. E.g Goch crucible and Sintered crucibles
Gooch crucibles
Sintered glass crucibles
3. Drying and ignition of precipitates
Drying and ignition are nearly the same except for the
temperature; drying is usually done at a temperature below
The crucible (or funnel) containing the precipitate is placed in
a small beaker and covered with a watch glass, and heated in
a drying oven at the prescribed temperature. In most cases,
this operation merely evaporates the water or other volatile
liquid, without causing any chemical change. A temperature
of 110o - 135o C is commonly used. When relatively unstable
substances (e.g. hydrates) are dried for weighing, the drying
temperature must be either low enough that no decomposition
occurs, or high enough to give complete conversion to a
stable decomposition product, (e.g. the anhydrous substance),
so that the composition of the weighing form is exact and
known. An initial drying period of 1 - 2 hours is used, then the
precipitate is cooled in a desiccators and weighed.
Heating for another half hour, cooling and weighing are
repeated till constant weight is obtained.
A precipitate is only dried if it has a definite chemical
composition and is only heated to remove its water of
hydration of crystallisation.
Ignition at temperatures above 250oC is done in order to
transfer a substance to a definite chemical composition; for
example precipitated Fe(OH)3 is first dried at 100oC to get rid of
the water adsorbed on the precipitate, then it is ignited to
remove the elements of water leaving the pure dry Fe2O3.
According to the nature of the precipitate, it may be
incinerated together or apart from the filter paper:
1. Incineration of the precipitate apart from the filter paper is
done for precipitates which are easily reducible by the
action of carbon produced on burning the filter paper, such
as copper oxide which is reduced with carbon to metallic
copper and BaSO4 which is reduced to BaS or Fe2O3 to Fe3O4,
etc. For this the precipitate after drying in the oven at 100o 105oC, is separated from the filter paper by friction, the filter
paper is first incinerated in a crucible until its carbon is
completely burnt off and the precipitate then added to the
crucible and incineration is completed until the crucible
acquires a constant weight.
Incineration of the precipitate and filter paper together: On the
contrary of the previous case, some precipitates are very stable
and unreducible by the carbon of the burnt filter paper, e.g.
Al2O3. If this is the case, the filter paper and the precipitate are
dried at 100o - 105oC, then the filter paper is folded down over
the precipitate to enclose it completely, and the whole is
placed in a tared crucible with the side having the three
thickness uppermost. The whole are pressed to the bottom of
the crucible without breaking the filter paper. The crucible is
then placed in an inclined position on a triangle, the cover
resting partly on the latter, and a small flame applied to slowly
cabonise the filter paper. The evolved gases should not be
allowed to catch fire, nor the paper, as this may causes
mechanical expulsion of the fine particles of the precipitate
owing to the rapid escape of the products of combustion. If
they inflame, the crucible should be covered to put out the
flame. When the paper is completely charred and no gases
evolve, the crucible is heated to red hot at the bottom, the
flame being kept away from the crucible’s mouth otherwise air
currents set up might sweep out particles of light precipitates;
also the reducing gases may reduce the precipitate. When all
carbon has been oxidised, the flame is adjusted to give the
required temperature of ignition of the particular precipitate
and heating is continued to constant weight.
Pyrolysis Curves:
Many precipitates of analytical interest decompose on heating
to give stable weighable forms, which can be weighed while it
actually in a furnace. A special balance designed for the
purpose is called thermobalance; it is sensitive to 0.2 mg. and
can be measure the temperature of the furnace to within
about 1oC, between room temperature and 110oC.
The graph recording the weight of the precipitate against the
temperature is called “pyrolysis curve”. It is evident that the
precipitate should be ignited in a temperature range where
the curve is flat, i.e. where the weight is constant over a wide
temperature range.
The pyrolysis curves of calcium oxalate and magnesium
oxalate are particularly interesting.
The momohydrate CaC2O4, H2O is stable at 100oC, then loses
water up to about 226oC, forming CaC2O4 which is stable up to
398oC, and then the oxalate loses carbon monoxide abruptly
to form CaCO3. The carbonate is stable in the range of about
commences. The weight finally becomes constant at about
850oC. Magnesium oxalate differs in its behavior in that it loses
magnesium oxide directly with no intermediate carbonate.
Burners of various types are used for burning or igniting the
precipitates. The more usually used is the blowpipe in which a
strong supply of air, in addition to the gas, in admitted from an
external air supply under pressure; the temperature available
by a blowpipe is from 1000 - 1200oC.
A sharp temperature for burning (400 - 1000oC) may be
arrived at by using an electrically heated “muffle furnace”
Desicators are intended for two purposes in gravimetric
analysis: storage of samples from which moisture has been
removed, and provision of a dry place in which samples, and
especially crucibles, may be to cool to room temperature after
heating or ignition and prior to weighing.
Weighing is the most important operation in both gravimetric
and volumetric analyses. In the latter, standard solutions are
prepared by dissolving a known weight in a known volume. For
gravimetric analysis the saught for material, or a derivative
thereof, is isolated from a known weight of sample and then
the pure material is weighed.
The analytical balance, is probably the most important tool in
reproducible. A part from the usual analytical balances,
damped balances are now commonly used.
A balance is said to be damped if it is subjected to some force
which restrains its motion and decreases the frequency of its
swings. Damping is induced by attaching to the beam either
large light pistons which move loosely in larger cylinder, or a
small aluminium plate which mover between the poles of a
powerful magnet. In the first case damping is due to friction of
the piston with the air inside the cylinder and the balance is,
therefore, called “Air-Damped”. In the other case, damping is
induced magnetically and the balance is the so-called
“Magnetically Damped”.
Calculation in gravimetric analyses:
In many gravimetric analyses the sought for material is isolated
free from other admixed ingredients and then weighed. From
the weight so obtained the percentage of the material can be
established. It is more common, however, not to isolate the
saught for material itself, but a compound or a derivative
containing it which is stable, easily weighable and having a
definite composition. Here the calculation of the amount of the
saught for material needs the application of our general
knowledge of the reactions taking place and the laws of
chemical combination.
For example, for the estimation of sodium sulphate in a sample
of impure salt or in a medicament etc., an aliquot is weighed
out, dissolved in water and the sulphate precipitated as barium
sulphate which is filtered off,
ignited and weight. Now the
reactions involved are :
Na2SO4 + BaCl2 =
BaSO4 + 2 NaCl.
BaSO4 is not altered by ignition.
That is, one molecule of barium sulphate is obtained from one
molecule of sodium sulphate, then the weight of sodium
sulphate present in sample can be calculated from the
wt.Na2 SO4
Mol.wt. BaSo4 Mol.wt. Na2 So 4
or Wt. Na2SO4 = Wt. BaSO4 x
wt. Na2 SO 4
Mol.wt. BaSo4
The ratio of the molecular weights of the materials involved, i.e.
wt. Na2 SO 4
, is called the gravimetric conversion factor.
Mol.wt. BaSo4
If, instead of sodium sulphate, aluminium sulphate is the
material to be estimated, then the reaction will be:
Al2(SO4)3 + 3 BaSO4 = 3 BaSO4 + 2 AlCl3
That is, one molecule of aluminium sulphate gives three of
barium sulphate and the conversion factor will be:
wt. Al2 SO4
3 x Mol.wt. BaSo4
It can be concluded that the molecular weight of the
substance sought appear in the numerator of the substance
weighed appears in the denumerator each multiplied by the
number of its molecules involved in the reactions.
If in the above examples the amount of an element only is to
be determined such as Na, Fe or S the atomic weight of such
elements replaces the molecular weight of the sought
compound in the numerator. The factor for sulphur in the
above estimation will be
At. wt. S
Mol.wt. BaSo4
Gravimetric conversion factors have been calculated for most
of the common analyses and extensive tables may be found in
most had books of chemical analysis.
Contamination of gravimetric precipitates
pharmaceutical analytical chemist in employing precipitation
as a means of separation and gravimetric determination is
obtaining the precipitate in a high degree of purity. However, a
precipitate formed in solution by the combination of suitable
reagents is almost invariably contaminated by the other ions
present during precipitation. It is necessary to investigate the
ways in which a precipitate can become contaminated and
what conditions can be employed to minimize contamination
and the methods can be used to increase the purity of the
precipitate after precipitation has been effected.
Contamination may be divided into two classes, according to
the mechanisms by which they associate with the precipitate.
A) Contamination by true precipitation
1. Simultaneous precipitation:
This occurs when the solubility product of the impurity is
exceeded during the addition of the precipitating agent, so
that both the impurity and the desired precipitate deposit
about the same time, for example, when silver chloride is
formed from a solution containing bromide ion as impurity, both
silver chloride and silver bromide precipitate.
importance as a source of error in gravimetric analysis. It is
possible to predict its occurrence by a knowledge of the
composition of the solution in which precipitation is carried out.
Moreover, it is possible to avoid such contamination, by
complexing. But the common methods of washing, digestion
and reprecipitation are seldom effective.
2. Conatmination by postprecipitation
Here, the contamination is introduced by the precipitation of
the impurity, sometime after the precipitate containing the
desired constituent has formed. This delayed precipitation is not
attributed to any great difference in solubilities between the
precipitate and impurity, but rather to a difference in rates of
crystallization. Moreover, postprecipitation is closely associated
with surface adsorption. The Ks for calcium oxalate is 2.6 x 10-9
and Ks for magnesium oxalate is 1 x 10-8, yet the latter does not
form in the absence of the former.
This suggests that the
surface of the calcium oxalate must be a major factors,
probably in the following way. Oxalate ion is present in excess
in the solution and therefore, comprises the adsorbed ion layer.
This effectively produces a relatively high concentration of
oxalate ion localized on the calcium oxalate surface to the
extent of providing a local state of supersaturation with respect
to magnesium oxalate, so that precipitation of the latter
Contamination by postprecipitation is apparently increased by
digestion, so that it is necessary to filter the saught for
precipitate as soon as its precipitation is complete, without
leaving it in contact with the mother liquor. In certain cases, a
water-miscible liquid may be added as soon as the primary
precipitation is complete in order to coat the precipitate
particles in such a way that their surfaces are no longer in
direct contact with the mother liquor. The organic material
used for coating must, of course, be volatile upon subsequent
ignition of the precipitate.
B) Contamination by coprecipitation
Some of the soluble materials contaminate the precipitate
and, if they cannot be removed by washing, are said to be
recognized; which of these predominates in a given case,
depends upon the particle size of the precipitate as well as the
chemical composition of both the precipitate and the solution.
The demarcation among these type is not sharp and,
furthermore, any
real case of coprecipitation may involve
more than one type.
1. Coprecipitation by isomorphic inclusion:
Two compounds are said to be “isomorphic” if they possess:
1. The same type of formula.
2. The same type of crystal structure.
3. The same type of bonding, and
4. Almost the same type of crystal lattice spacing.
Thus, if the precipitate and an impurity are isomorphic, one ion
within the crystal of the precipitate is replaced (during the
precipitation) right in the crystal lattice by an ion of the impurity
which actually becomes incorporated permanently into the
crystal lattice of the precipitate. The impurity is distributed
throughout the host crystal of the precipitate, sometime
uniformly, and sometimes non-uniformly, for example, BaSO4
and PbSO4 , are isomorphic, they have the same type of
formula, the same type of crystal structure, the same type of
bonding (both divalent salts) and almost the same type of
crystal lattice spacing (ionic diameters of Ba2+ and Pb2+ are 5 Å,
respectively (the angstrom, Å is 10-8 cm). If SO4- is added to a
solution containing Ba2+ and traces of Pb2+ in concentrations
not sufficient to exceed the solubility product of PbSO4,
contamination of the precipitate with lead sulphate occurs
despite the fact that the solubility product of the lead
compound is not exceeded. The lead ions replace some of the
barium ions in a random fashion throughout the crystal lattice,
causing no appreciable distortion of the barium sulphate
This coprecipitation, resulting in mixed crystal formation,
cannot be appreciably diminished by the ordinary methods of
reducing coprecipitation, viz., digestion, and washing. The only
real way to eliminate this type of coprecipitation error, is to
remove the contaminating ion beforehand, or to dissolve the
contaminated precipitate and to carry out reprecipitation
under more favorable conditions.
2. Coprecipitation by non-isomorphic inclusion.
When two substances are not completely isomorphic to be
miscible in all proportions, they may still be soluble to some
extent in each other. Such may be the case when two
compounds have the same type of formula, the same crystal
structure; but different bonding and different crystal lattice
spacing. This may be illustrated by the contamination of barium
sulphate with potassium ion. The first is a divalent salt, while, the
latter is a monovalent salt. the ionic diameters for Ba2+ and K+
are 5Å and 3Å, respectively. Barium sulphate may to lerate
limited amounts of potassium without sufficient distortion to lose
its characteristic crystal symmetry and structure. In small
amounts, the contaminate may be distributed throughout the
host in a manner similar to isomorphic inclusion. It should be
noted, however, that the replacement of one barium ion with
one potassium ion must be complained by some other
modification of the crystal lattice in order to maintain electrical
neutrality. Space considerations would certainly not permit two
potassium ions to occupy the lattice site of a single barium ion.
However, a hydrogen sulphate ion, HSO4-, is similar to a
sulphate ion, so electrical neutrality is preserved if a hydrogen
sulphate anion replaces a sulphate ion every time a potassium
ion replaces a barium ion. That this must occur, is indicated by
the observation that coprecipitation of potassium ion is almost
negligible where very few hydrogen sulphate ions are present
(at pH 5); but coprecipitation of potassium ions is appreciable
at pH 1, where that of the sulphate exist as hydrogen sulphate.
3. Coprecipitation by surface adsorption:
As already described ions are adsorbed from the mother liquor
into the surface of precipitated particles. This adsorption
involves a primary adsorbed-ion layer held very tightly, and a
counter-ion layer held more or less loosely. These ions are
carried down and distributed over the particle surfaces of the
precipitate. As the case with gelatinous precipitate and, to
some extent with curdy precipitates, it is seldom important for
coarsely crystalline precipitates. Where surface adsorption is
the predominant mechanism, the bulk of impurity may be
removed by simple washing.
4. Coprecipitation by occlusion:
An adsorbed, nonisomorphic impurity never fits as well into a
crystal structure as the host substance itself. Hence even
through the impurity may be temporality adsorbed on the
surface of a crystal, it tends to be pushed out and replaced by
the host substance as the crystal grows, so that the interior of
the crystal is uncontaminated the impurity long restricted only
to the surface of the particle. However, if the crystal grows
rapidly, the impurity does not have time to be deebed, it may
be covered over and entrapped within the rapidly growing
crystal. Such enclosed impurities are said to be occluded or
internally adsorbed. This type of impurity differ from isomorphic
1. The extent of occlusion may not follow any particular
laws, but may be very variable and highly dependent
upon the mode of formation of the precipitate.
2. The
distributed throughout the precipitate.
3. It is possible to remove the occluded impurity by
washing, if the wash liquid can reach it. Hence digestion
and reprecipitation are often effective in removing
occluded impurities.
In another mode of occlusion, impurities are entrapped in the
spaces between the particles in aggregates. These voids may
trap not only impurities adsorbed on the surface of the void,
but also mother liquor with its impurities, whether they are
adsorbable or not.
Reduction of contamination
Measures taken to diminish the contamination, and hence the
error, include:
 elimination of contamination
 digestion of the precipitate
 reprecipitation and
 Washing.
1. Elimination of contamination
The contamination may be eliminated by actual separation
before precipitation; or by changing it chemically to a less
objectionable form without actual separation, by employing
complxing reactions, redox reactions, ...etc. For example,
ferrous iron is adsorbed much less than ferric iron to a
precipitate of barium sulphate. Therefore, in the determination
of barium by precipitation as barium sulphate, any iron present
should be reduced before to the ferrous condition.
2. Digestion:
Digestion, i.e. allowing the precipitate to stand in contact with
the mother liquor for some time before filtration, is often an
effective method in lowering the degree of contamination of
crystalline precipitates. Several processes take place during
A freshly-formed precipitate contains micro and macro
crystals, the former are usually more soluble than the latter. This
render the mother liquor supersaturated with respect to the
larger crystals, and in order to establish equilibrium, additional
material leave solution and enter the solid phase depositing on
the large crystals, causing them to grow even larger. The
impurities which were held by the micro crystals are not
adsorbed appreciably by the large crystals. This process is
called Ostwald ripening process. The net result is the
disappearance of small impure crystals, and the growth of
purer and more perfect larger crystals.
Another process, internal ripening, brings about an internal
perfection of the precipitate crystals. Thus, a freshly form
precipitate has a rather irregular form and contains many
imperfections. On standing in contact with its mother liquor,
ions from the lattice enter the body of solution and are
replaced by ions from the solution. This is attributed to the fact
that ions form the protruding point at the corners and edges of
the crystal lattice dissolve more readily than others, while ions
replacing them enter the recesses and ore vices in the lattice in
an orderly fashion. The net result is a general perfection of the
crystal lattice. During this time, occluded and adsorbed
impurities may be expelled, and since the number of surface
ions has decreased, the amount of readsorbed impurities
decreases also, and a general purification of the precipitates.
For coagulation colloidal precipitates, digestion -as a
means of purification- is of little value due to the very large
Moreover, the solubility of soft substances of low surface tension
is independent of the particle size. Therefore the micro crystals
will not dissolve to be replaced on larger crystals, leaving back
their impurities. However, for some curdy precipitates, such as
silver chloride, digestion causes a decrease in the total surface
area and a corresponding decrease in the amount of
contamination, but with no growth in size; except that the tiny
colloidal particles may form loose agglomerates by sharing
their water jackets.
For gelatinous precipitates, however, digestion is not
recommenced, it may render the precipitate slimy and thus
impairs its filterability.
The stable form of the precipitate may be attained
during digestion. Often, the precipitate first formed may be
an allotropic modification of the saught- for precipitate. Such
forms are often amorphous, and consequently impure or
unfilterable. Being less stable, however, they have higher
solubilities than the stable form during digestion. In the
recrystallisation process, impurities are returned to the mother
liquor, and the stable form is usually purer.
In a reprecipitation, a sought-for
precipitate is washed, dissolved in pure solution and
repreprcipitated. This method is satisfactory in decreasing the
surface-adsorbed and occluded impurities, it may also
postprecipitation. During the second precipitation, the
amount of contaminant adsorbed will be lower since its
concentration in the solution is less. As a result, less will be
enclosed within the crystalline structure of the precipitate
when it is again formed. But reprecipitation is of little value in
precipitation or by isomorphic inclusion.
Reprecipitation is
also unsatisfactory whenever the contamination arises from
the reagent used for precipitation process itself.
Thus, contamination of hydroxides with basic salts formed
during the precipitation process cannot be greatly reduced just
by reprecipitation.
Because reprecipitation requires much time and extra
manipulation which increase the probability of error, and
because the procedure is in application wherever no simple
methods of dissolution are available, reprecipitation is not
generally desirable.
Washing. Washing a precipitate serves to remove
surface adsorbed impurities, and impurities in the mother
liquor held mechanically by the precipitate. Simple washing is
precautions must be taken in choosing a wash liquid :
1. It should not dissolve any considerable amount of the
2. It should not peptize the precipitate.
3. It should neither metathesize the precipitate nor cause
MgNH4PO4 tends to hydrolyze to MgHPO4 when washed
with water, and must be washed with dilute NH3 to
prevent such hydrolysis.
Hydrolysable metal salts, such as those of Fe3+ and Al3+,
cannot simply be washed away from a sought-for precipitate
precipitate. Dilute nitric acid may be used to prevent such
Using several small portions of wash liquid are more effective
than washing with few larger portions. Washing crystalline
precipitate may be done with water, and if the precipitate is
somewhat soluble, a saturated solution of the precipitate may
be used as the wash liquid. Washing can also done with
organic solvent, like alcohol and ether, if the precipitate can
be dried at a low temperature.
Coagulated colloidal precipitates require care in washing.
Thus curdy precipitates with their large surface area are more
liable to contamination than crystalline precipitates, and are
therefore prone to peptization if washed with water. For this
purpose, a solution of a strong electrolyte such as nitric acid is
used for washing. The ions of the electrolyte replace the
adsorbed contaminant ions, and at the same time prevent
peptization of the precipitate. The electrolyte chosen should be
volatile so that it will be removed during subsequent drying or
ignition of the precipitate. In case of gelatinous precipitates,
made up of very small particles with huge surface area, the
changes on the particles are controlled by the pH of the wash
solution; and are usually ashed with an electrolyte solution such
as ammonium nitrate. This prevents peptization of the
precipitate, and the ions adsorbed are largely replaced by the
volatile ions of ammonium nitrate.
Organic precipitants
Certain types of organic compounds produce precipitates
with metal ions with a high degree of selectivity. The
precipitates formed are usually crystalline and pure, they are
easily filtered and washed, and contain a low metal ion to
precipitate ratio. Furthermore, their selective action can often
be improved by pH control and the use of masking. Most
organic precipitants combine with cations to form chelate rings
(non-ionic complexes). Other organic precipitants form salt-like
precipitates with metal ions.
1. Precipitants forming chelate compounds:
compounds with cations, contain both an acidic and a basic
coordinating (electron-donating) functional group. The metal,
interacting with both of these groups, becomes itself one
member of a heterocyclic ring. From the strain theory of
organic chemistry, it is expected that rings of this type would be
mainly five- and six- membered. Hence, the acidic and basic
functional groups in the organic molecule must be situated in
positions with respect to each other which permit the closure of
such rings. Insofar as the chalets are non-polar, they will have
low solubility in water and high solubility in organic solvents.
Chelating precipitants are weak acids: therefore in forming the
insoluble complex salts, the usual equilibrium principles of
ionisation constants and solubility product constants apply, for
example, if:
H.Chel represents the organic precipitant.
represents a cation of valence number “n”
 H+ + Chel-
Ka =
[H ][Chel ]
and K a 
n.Chel- + Mn+  M [Chel]n
and Ks = [Chel-]n [Mn+]
The overall reaction can be represented thus:
n H.Chel + Mn+
 nH+ + M(Chel)n
A neutral chelate compound of the type M(Chel)n is essentially
organic in nature. The metal ion becomes simply one of the
members of an organic ring structure, and its usual properties
and reactions are no longer readily demonstrable.
The following are some of the acidic, and coordinating groups
commonly found in organic chelating agents; note that the
oxime group appears in both categories.
Acidic groups
Coordinating groups
- COOH (carboxyl)
-NH2 (amino)
-OH (hydroxyl)
-NO (nitroso)
-NH (imine)
=N- (cyclic nitrogen)
= N-OH (oxime)
= N-OH (oxime)
-SO3H (sulphonic)
= O (carbonyl oxygen)
In order to form a chelate with a given organic precipitant,
the metal ion must be of appropriate size, oxidation number,
and coordination number. In many cases, the chelate
compounds formed are suitable weighing forms for gravimetric
analysis. When this is the case, the gravimetric factor for
calculation to the metal is very favorable, because the metal
constitutes only a small part of the heavy chelate molecule. If
the precipitate is not a suitable weighing form, it can be ignited
to the metal oxide for weighing.
2. Precipitants forming salt-like compounds:
produce slightly soluble salt-like
compounds (rather than chalets)
in which the bond with the species
precipitated is primarily ionic in
Sodium tetraphenylboron:
This compound has the formula Na+B(C6H5)4- and is a specific
precipitating agent for potassium and ammonium ions in cold
mineral acid solution. Under these conditions, only mercuric
mercury, rubidium, and cesium interfere, and must be removed
before treatment.
This compound is used primarily for the precipitation of
sulphate ion from a slightly acid medium.
The solubility of the precipitate increases rapidly with
temperature and with the acidity of the environment; both of
these variables must be carefully controlled. The precipitate is a
satisfactory weighing form.
Substituted arsonic acids
Substituted arsonic acids have the structure.
Where, R is an organic group, especially
phenyl, p-hydroxyphenyl, and n-propyl. These
acids precipitate quadrivalent metal ions such
as tin, thorium, and zirconium from acid media.
The precipitate generally contains two moles of the acid per
mole of quadrivalent cation. The nature of the organic portion
of the molecule (R) determines, to some extent, the cations
that form precipitates as well as the conditions under which
they are formed. Because these precipitates are difficult to dry
without decomposition, metallic aresonates are ignited to their
respective oxides before weighing.
Application of gravimetric analysis
An important consideration in gravimetric analysis is the
specificity of the reagent used for the formation of the
precipitate. Ideally, a single element in the periodic table
would precipitate with a given reagent under a given set of the
reaction conditions. In fact however, very few of the reactions
available even approach such specific
behavior. Most
reagents, at best, are selective, i.e. they precipitate several ions
at the same time. In some cases, improved selectivity can be
obtained by oxidation or reduction of some of the ions present,
by control of the pH of the solution, or by the use of masking
agents. For example, the precipitation of chromic chromium
along with ferric iron as hydroxides on the addition of
ammonium hydroxide does not occur if the chromium is first
oxidized to chromate. The use of pH control to increase the
selectivity of metal sulphides and metal hydroxide precipitation
is well known. Similarly, the precipitation of aluminium along
with ferric iron as hydroxides can be prevented by increasing
the pH of the solution to form aluminate. So also masking
agents may be used to inhibit certain precipitations and
enhance selectivity. Familiar examples are the use of tartaric
acid to prevent the precipitation of iron as iron hydroxide and
prevent precipitation of copper as hydroxide (in Fehling
solution); and the use of ammonia to prevent the precipitation
of chloride as silver chloride.
Selected gravimetric methods
Several specific determinations, in each of which the final
quantitative measurement is that of weight, are discussed in
this chapter. One involves a quantitative separation by
volatilisation, and each of the other includes a quantitative
selected to encompass a variety of precipitates, procedures,
and principles. Considerable emphasis is placed on a
description sources of error; these are discussed in terms of
the following requirements:
1. The
2. The precipitate must be quantitatively pure or of
known purity at the time of final measurement, and
3. The precipitate must be in a physical form suitable for
subsequent handing.
1. Determination of chloride
The determination of chloride as silver chloride involves all those
operations necessary for precipitating a fairly insoluble, curdy
precipitate which is easily washed free of impurities.
The chloride content of a soluble salt can be determined by
precipitation as silver salt.
Ag+ + Cl-  AgCl
The precipitate is collected in a filtering crucible, washed, and
brought to constant weight by drying at 110oC.
Completeness of precipitation:
The solubility of silver chloride increases on heating, therefore
the precipitate should always be washed and filtered at room
temperature. A slight excess of silver nitrate should be added to
reduce the solubility of silver chloride (common ion effect). A
dilute nitric acid medium is required to hasten and to maintain
flocculation any solubility loses. Cations which form soluble
chloro- complexes, e.g. mercuric ions which give soluble HgCl42, should be removed prior to precipitation.
Purity of the precipitate;
The specificity of the precipitating agent for chloride ion is
much improved in the presence of nitric acid. Most of the
anions of weak acids, e.g. phosphate, carbonate, ..etc., which
form insoluble silver salts in neutral medium only, will not,
interfere. However, even in dilute nitric acid, bromide, iodide,
and thiocyanate ions still yield insoluble precipitates with silver
ions, and these ions must be removed beforehand. Cations
which hydrolyse on boiling the solution to give insoluble
oxychlorides, e.g. bismuth, antimony, and ferric ions:
Bi3+ + H2O
Cl  2 H+
+ BiOCl
must be removed in advance. It is apparent from the above
hydrolytic equilibrium equation that hydrogen ion (in the form
of nitric acid in this case) would displace the eqilibrium system
by dissolving the insoluble oxychloride. But this needs high
concentration of acid and would introduce error through
oxidation of the chloride ion:
2Cl- + 2 NO3- + 4 H+  Cl2 + 2 NO2 + 2 H2O
Silver halides are decomposed into their constituent elements
on exposure to light specially during the early stages of the
precipitation when the milky suspension presents an enormous
surface for photochemical reaction. The mechanism of this
decomposition is apparently complex, but may be represented
for simplicity as follows:
2 AgCl  2 Ag + Cl2
accumulation of finely divided silver.
Agents which reduce silver salts into metallic silver, e.g.
formation, sulphur dioxide, reducing ions, etc., must be absent.
The precipitate should be filtered through filter crucibles. If filter
paper is used, it will largely reduce the precipitate adhering to
it during heating in the drying operation. In such a case, the
filter paper is burnt, the ash treated with nitric acid to dissolve
any reduced silver, then hydrochloric acid is added to convert
it back to silver chloride.
Washing and drying:
To suppress peptization, the precipitate is washed with a
volatilisable electrolyte, e.g. very dilute nitric acid. the
precipitate may finally be washed with ethanol to remove
much of the water and, being more volatile than water, assists
in the drying process. However, the precipitate should not be
dried rapidly, otherwise it would shrink and trap water that
cannot be completely removed except on heating to melting
point of the precipitate, but then decomposition begins.
2. Determination of Iron
The gravimetric determination of iron is almost invariably made
by oxidising it (if necessary) to the ferric state, and precipitating
it with ammonia as ferric hydroxide, filtered on paper, washed
again dried, and ignited to the oxide and weighed:
2 Fe3+ + 6 NH3 + (x+3) H2O  Fe2O3 . x H2O
+ 6NH4+
Fe2O3 . x H2O  Fe2O3 + H2O
Completeness of precipitation:
The precipitation of iron as hydrous ferric oxide is not complete
if the iron is not entirely in the ferric state at the time of
precipitation. Any ferrous iron must be oxidised beforehand, by
means of bromine, or nitric acid, or hydrogen peroxide.
+ Br2
 2 Fe3+ + 2 Br-
3 Fe2+ + NO3- + 4 H+  3 Fe3+ + NO + 2 H2O
2 Fe2+ + H2O2 + 2 H+  2Fe3+ + 2 H2O
The latter oxidant is preferable, because its reduction
product is water, and because any excess peroxide easily
decomposes by boiling the solution briefly.
2 H2O2  2 H2O + O2
The primary particles of the precipitate are extremely minute,
and their enormous surface development is conductive to
extensive contamination by adsorption.
Purity of the precipitate:
Nearly every metallic ion other than alkali metal ions can form
a hydroxide, hydrous oxide, or basic salt in alkaline solution,
and so all may interfere and impair the purity of the precipitate
of ferric hydroxide. Careful control of the hydroxyl ion
concentration is necessary to provide the required selectivity.
Buffers , as well as weak bases (as the may be ) are to be used
in limiting the hydroxyl ion concentration to the predetermined
value; through complications often ensure by momentary local
excesses of hydroxyl ions where the drops of added reagent
enter the solution inn the precipitation vessel. This particular
source of difficulty is eliminated in the method of precipitation
from homogeneous solution.
Dispositive ions such as Zn3+ and Cu2+ which form soluble amino
complexes, nevertheless, cause serious interference. Thus
concentration is low, the hydroxide, which obstructs them from
the solvent action of excess ammonia, and even if these
complexation (M(NH3)4)2+ are quite strongly adsorbed on and
carried down with the iron hydroxide in alkaline solution.
Tri- and tetra-positive cations such as Al3+, Cr3+ and Mn4+ are
precipitated as their hydrous oxides along with the ferric iron.
Although Mn2+ is not quantitatively precipitated by ammonia in
presence of high concentration of ammonium ion, yet Mn2+ in
the original solution is oxidised along with ferrous iron if bromine
is used as the oxidising agent, and is quite easily air-oxidised
and precipitated as Mn4+ hydrous oxide in the alkaline solution.
therefore, these cations must be removed before precipitating
iron in the course of its gravimetric determination.
Repecipitation is often necessary in order to obtain proper
precipitate of ferric hydroxide. But if the precipitate is highly
impure, it is preferable to use a fairly specific or organic
Filtration and washing:
Ferric hydroxide as all gelatinous precipitates, retains within the
solid phase a quantity of the mother liquor. The precipitate
tends to clog the pores of the filter and therefore filter crucibles
cannot be used for filtration. A coarse filter paper should be
employed, and nearly all of the washing operation should be
accomplished in the precipitation vessel by decantation prior
to transfer of the main body of precipitate to the filter.
Filter paper can be added to the solution and upon which the
precipitate gathers, as a means to effectively increase the
bulkiness and consequently the ease of handling this gelatinous
As wash liquid, a hot dilute solution of an electrolyte such as
ammonium nitrate is generally used, the electrolyte serving to
prevent peptization. After transferring the precipitate to the
filter, washing must be continued without interruption. If the
gelatinous precipitate is allowed to dry, it shrinks and cracks in
the paper, and any wash liquid poured on such a precipitate
runs through the channels and does not accomplish much in
washing out the soluble impurities.
Upon drying the precipitate, a high temperature must be used
to remove all traces of water. A final ignition temperature of
about 1000oC. is suitable. If this temperature is exceeded, some
of the ferric oxide may be reduced to the magnetic oxide by
loss of oxygen.
red .
6Fe2O3  O2 + 4 Fe3O4
Reduction of the ferric oxide to magnetic iron or even to
metallic iron may occur by the carbon of the filter paper, or by
the reducing gases from the flame if allowed to envelop the
crucible containing the precipitate during ignition.
Another source of error occurs if the precipitate contains much
ammonium chloride if the latter has been used incorrectly as
wash liquid, whereby ferric chloride may form and volatilise
during ignition;
Fe(OH)2 + 3 NH4Cl  3 NH2 + 3 H2O + FeCl3
But if ammonium nitrate has been used as the wash liquid, no
such error occurs, because the ferric nitrate formed will
decompose on ignition giving ferric oxide:
4 Fe(NO2)3  12 NO2 + 2 Fe2O3 + 3 O2
3. Determination of Aluminium.
This is another gelatinous precipitate; aluminium is precipitated
as its hydrous oxide and ignited to the oxide. It is subject to the
same coprecipitation errors that were discussed under the
precipitation of iron. However, the separation of aluminium is
amphoteric which limits the pH range over which it may be
hydroxide begins at pH of about 3 and is complete near the
neutral point,, but is re-dissolve at a pH of about 9.
Aluminium hydroxide ionises as either a base or an acid:
Al(OH)3  Al3+ + 3OHAl(OH)3  H+ + H2AlO3Consequently, aluminium hydroxide dissolves in strong acid,
and in excess alkali. There is some pH at which aluminium
solubility; this isoelectric point, as calculated from the above
two equations, is pH 5.3.
Two washed precipitate is dried, and ignited briefly above
1200oC to expel moisture completely, and to convert it to a
different crystal modification which is not hygroscopic.
Aluminium may also be quantitatively precipitated by oxime
(8-hydroxy quinoline) from an acetic acid-acetate buffer
solution. The precipitate is crystalline, little contaminated by
adsorption of other ions from solution, and is easy to filter and
wash. The precipitate is suitable form weighing as such, after
oven-drying at about 130oC.
If the aluminium solution contains cations that precipitate
under the condition of experiment, it is advisable to separate
the aluminum of the experiment, it is advisable to separate the
aluminum by precipitation. Thus, if cations of nickel, cobalt,
manganese, zinc, magnesium, etc. are present, the solution is
first neutralised, and then reacted with ammonium acetate
and boiled. Aluminum precipitates as the basic acetate,
Al(OH)2(CH2COO), which is filtered, washed, dissolved in acid,
and then precipitated from the now pure solution as hydroxide.
On the other hand, if the interfering ions are alkaline earth ions,
the aluminum solution is treated with acetic acid and
acid in order
the aluminum
phosphate AlPO3. This is filtered, washed, dissolved in acid and
precipitated as above.
4. Determination of sulphate.
Sulphate ion solution is determined by precipitation with
barium chloride from a hot solution slightly acidified with
hydrochloric acid. The precipitate is digested, washed, ignited
and weighed. The determination centers on the reaction:
Ba2+ + SO42-  BaSO4
Barium sulphate is classified as a crystalline precipitate, the
individual crystals are much larger than they are in a curdy
precipitate such as silver chloride.
Completeness of precipitation.
The solubility of barium sulphate at room temperature (about
3 mg. per liter) is sharply diminished by the presence of excess
barium ion is the mother liquor; and is increased somewhat at
elevated temperatures.
Barium sulphate is markedly more soluble in acid media than in
pure water. The second ionisation of sulphuric acid is slight, K s =
3 x 10-2 ; hence hydrogen ion can lower the sulphate ion
concentration to some extent and thus increase the solubility of
barium sulphate:
BaSO4 + H+  Ba2+ + HSO4By precipitating barium sulphate in hot dilute solution and in
supersaturation is decreased somewhat, so that a better
precipitate is obtained., The acidity contributes to the purity of
the precipitate as will be shown.
Purity of the precipirtate:
Barium ion forms insoluble precipitates with a variety of anions
other than sulphate ion, including carbonate, phosphate,
precipitate with the sulphate. Yet most of them are anions of
weak acids so that their barium salts are soluble in acid media.
The only anion which remains insoluble under these conditions
(of acidity) is the fluoride ion, it is removed by volatilisation as
hydrofluoric acid, or by complexation with boric acid.
Some anions, such as nitrate and chlorate, are coprecipitated
as their barium salts to such an extent as to introduce very large
errors. Nitrate and/or chlorate in the original sample can be
removed by repeated evaporation with hydrochloric acid:
2 NO3- +
6 Cl- + 8 H+  2 NO + 3 Cl2 + 4 H2O
+ 5 Cl- + 6 H+  3 Cl2 + 3 H2O
The more common cations which simultaneously precipitate as
sulphate along with barium sulphate, are lead, calcium, and
strontium. Lead may be removed by complexation with
acetate, while calcium and strontium are removed by fusing
the sample with sodium carbonate. The melt is extracted with
water, and the calcium and strontium (now present as insoluble
carbonates) are removed by filtration. The filtarte, containing
the pure sulphate, is acidified boiled to remove carbon dioxide
prior to precipitation as barium sulphate.
Several heavy metals, whose sulphate are soluble, tend to
coprecipitate with barium sulphate; the higher the oxidation
number of the cation, the greater the coprecipitation. Metallic
reducing agents will reduce the oxidation number of many
cations such as ferric iron; and will also remove cations of
inactive elements such as copper, antimony by reducing them
to the elemental state. Alkali metals may also coprecipitate as
sulphate along with the barium sulphate.
Ammonium salts, if coprecipitated, cause serious error and
must be removed before precipitating barium sulphate. The
removal is done by evaporation with nitric acid:
NH4+ + NO3-  N2O + 2 H2O
hydrochloric acid.
Barium sulphates not sufficiently soluble in any reagent to
permit dissolution and reprecipitation as a means of decreasing
coprecipitated with barium sulphate and their influence on the
results of analysis are mentioned in the following.
The coprecipitation of anions as nitrate and chlorate leads to
positive error, i.e. high result. Thus, barium nitrate ignites to the
oxide, and barium chlorate ignites to chloride, and these
residues add to the weight of the barium sulphate.
The coprecipitation of cations leads to either positive or
negative errors. Thus, coprecipitation of ammonium sulphate
gives negative error, because of the complete volatilisation on
(NH4)2SO4  2 NH3 + SO3 + H2O
The coprecipitated sulphates of heavy metals undergo
decomposition during ignition, losing sulphur trioxide and
leaving a residue of the metal oxide. Whether the error is
negative or positive depends upon the relative weights of the
sulphur trioxide volatillised and of the metal oxide remaining as
residue. Of the cations commonly encountered ferric iron is
noteworthy. In this case, the precipitate is partially ferric
sulphate which is decomposed upon ignition to ferric oxide.
Each sulphate ion should account for one barium sulphate
molecule (molecular weight 233) in the final precipitate, but
since each sulphate ion involved in the coprecipitation of iron
accounts for only one
third of a ferric oxide molecule
(molecular weight 159.69, one third being
only 53.23), the
precipitate is too light, and the result for sulphate comes out
low; i.e., the error is negative. Coprecipitation of other cations
leads to a similar conclusion weighs less than the barium ion
itself, and vice versa.
Physical form of the precipitate:
Although barium sulphate is crystalline precipitate, yet it
exhibits a tendency to “creep”, i.e. the fine clumps of
precipitate, supported on a liquid surface and moving over it
through the action of surface tension, distribute over the entire
wetted surface of the containing vessel, even surfaces are wet.
Thus, the particles of the precipitate can climb up a wetted
filter paper into the side of the glass funnel, if the latter is wet,
and may be lost. In such cases, the filter paper must be filled to
the rim. Digestion helps decrease creeping of the precipitate.
To remove moisture completely from the washed dried barium
sulphate precipitate, it must be ignited at a temperature of
500oC. or above. The barium sulphate itself is stable well above
this temperature, but if oxidising conditions to hot prevail during
ignition, it may be reduced by the carbon of the filter paper or
by the carbon monoxide of the flame if the latter completely
envelops the precipitate:
BaSO4 + 2 C  BaS + 2 CO2
BaSO4 + 4 CO  BaS + 4 CO2
If reduction occurs, the results are too low, but the error can be
corrected either by heating with good access air:
BaS + 2 O2  BaSO4
or by moistening the residue with one or two drops of
concentrated sulphuric acid:
BaS + H2SO4  BaSO4 + H2S
Volatilisation of the excess sulpuric acid must be done very
carefully to avoid loss by spattering; evaporation over a free
flame will almost certainly result in mechanical loss, therefore
use of a radiator or air bath is required.
Almost any form of sulphur can be determined by first
converting it to sulphate and then precipitating as barium
sulphate. Sulphur, sulphides, and sulphites are
oxidised to
sulphate; while peroxydisulphates (persulphates) are reduced
to sulphate.
Of the official substances, cited in the Pharmacopoeias, the
following are worth mentioning. Paromomycine sulphate,
sodium sulphate and ichthammol. Ichthammol consists of the
ammonium salts of certain sulphonic acids together with
ammonium sulphate. It is analysed for both organically
combined sulphur and for sulphate:
a) For sulphate, the ichthammol sulphonic acids are removed
by shaking with cupric chloride solution and filtrating off the
precipitated cupric salts. The free sulphate, which is present
in the filtrate, is determined by precipitating as barium
b) For organically combined sulphur, the ichthammol is
heated with copper nitrate and sodium carbonate. The
mixture is acidified with hydrochloric acid to convert the
precipitated and determined as barium sulphate in the usual
manner. From the percentage of total sulphur thus obtained,
the percentage of sulphur in the form of sulphate is
subtracted to give the percentage of the organically
combined sulphur.
Barium, strontium and lead may be determined by adding an
excess sulphate ion, just the reverse of the sulphate procedure.
coprecipitation leads to low results, and cation coprecipitation
gives high results- just opposite to the situation in sulphate
determinations. The solubilities of strontium and lead sulphates
are too great for good quantitative work, but are rendered less
soluble in a mixed solvent of alcohol and water.
5. Determination of calcium.
Calcium ion is precipitated as the oxalate monohydrate and
weighed as such:
Ca2+ + C2O42- + H2O  Ca2C2O4. H2O
Other weighing forms are the carbonate, the oxide (obtained
by heating the oxalate to the proper temperature) or the
sulphate. The latter is obtained by treating the oxalate with
sulphuric acid, and igniting to convert it into calcium sulphate:
CaC2O4. H2O + H2SO4  CaSO4 + CO + CO2 + 2 H2O
Calcium oxalate is soluble in strong acids, due to lowering of
the oxalate ion concentration by formation of bioxalate ion,
HC2O4-. Calculations from Ksp of calcium oxalate and K2 of
oxalic acid show that the solubility of calcium oxalate in
solution of pH 5 or 4 is only slightly greater than in neutral
solution. The solubility in slightly acidic solution is decreased by
the common ion effect of excess precipitant.
Most of the common metals, except the alkali metals, form
sparingly soluble oxalates and must be removed before the
calcium is precipitated. Magnesium oxalate, however, is only
moderately soluble in water, and tends to form stable
supersaturated solutions. the solubility of magnesium oxalate is
increased by adding excess of oxalate ion, due to formation of
oxalate decreases the solubility of calcium oxalate, due to the
common ion effect.
Use in made of the above two properties of calcium oxalate in
the precipitation process to obtain crystalline, large particles
precipitate without digestion (which would increases the
danger of post precipitation of magnesium oxalate). Thus, the
solution of the sample of calcium salt is acidified with
hydrochloric acid, heated, and ammonium oxalate is added.
The acidity, which prevents the precipitation of calcium
oxalate, is neutralised homogeneously by adding an excess of
urea which undergoes slow hydrolysis, and the ammonia
liberated reacts with the bi-oxalate ion and thereby increases
the oxalate ion concentration and effecting precipitation of
the calcium oxalate.
The precipitate is digested for an hour, then filtered promptly.
Longer digestion may result in considerable contamination by
postprecipitation of magnesium oxalate. The precipitate is
washed with electrolyte solution, e.g. ammonium oxalate, then
dried at about 100 oC., to remove the moisture. Then the filter
paper is charred, and ignited till all the carbon has been burnt
off. The cold residue is moistened with a mixture of sulphuric
acid and ethanol, heated, to remove the acid; and the
process is repeated until constant weight is attained.
6. Determination of magnesium
The determination depends upon the addition of phosphate
ions and ammonia to precipitate the magnesium ions as
crystalline hydrated magnesium ammonium phosphate which
is then ignited to the pyrophosphate and weighed:
Mg2+ + NH4+ + PO43- + 6 H2O  MgNH4PO4. 6H2O
2MgNH4PO4. 6H2O  Mg2P2O7 + 2NH3 + 13 H2O
precipitate is fairly soluble in water, but much less so in the
presence of ammonia. It forms slowly and therefore the
precipitate must be allowed to stand for several hours before
filtration. There is a tendency for the precipitate to be
contaminated with other phosphates if much potassium,
sodium, or ammonia ions are present in solution, e.g. MgKPO4,
MgNaPO4, and Mg3(PO4)2 all of which are unchanged on
ignition thus leading to negative error; and Mg(NH4)4(PO4)2
which gives magnesium metaphosphate on ignition leading to
high results:
Mg(NH4)4(PO4)2  Mg(PO3)2 + 4 NH3 + 2 H2O
Therefore, the excess ammonium salts must be removed
before precipitation of the magnesium. Reprecipitation may
also be resorted to, especially if excess potassium and sodium
ions are also present.
The magnesium ammonium phosphate is allowed to stand with
the mother liquor several hours in the cold, filtered and washed
with cold dilute ammonium hydroxide to minimise solubility loss,
and dried. The filter paper is charred, and burned off at the
lowest possible temperature until all carbon is destroyed. The
precipitate is then ignited at about 1000oC, to convert it to the
weighing form; magnesium pyrophosphate. Prolonged heating
at 1200oC results in slow loss of phosphorus pentoxide. The
ignited precipitate is not hygroscopic, and offers no difficulty in
Magnesium may also be determined by precipitation with
oxime (8-hydroxyquinoline). The precipitate forms rapidly in
ammonical medium; contamination by coprecipitation of
other inorganic components of the solution is negligible,
through some of the reagent may be coprecipitated. Weighing
can be made of the dihydrate, Mg(C3H5OH)2. 2H2O, after ovendrying at 105oC. It is usually better to dry at about 140oC to
remove all the water of hydration as well as any coprecipitated
7. Analysis of mixture
A) Analysis of mixture of magnesium and calcium:
The calcium is precipitated as oxalate, using excess oxalate
to keep the magnesium in solution. The calcium oxalate is
filtered rapidly, reprecipitated, and determined as explained.
The filtrate, containing the magnesium oxalate complex, is
treated with nitric acid to oxidise the excess oxalate:
C2O42- + 2 NO3- + 4 H+  2 CO2 + 2 H2O + 2 NO2
The magnesium is then determined as the pyroposphate, or
as the oxinate.
B) Analysis of mixture of aluminum and magnesium
Aluminium can be precipitated with 8-hydroxyquinoline from
acetic acid- acetate buffer mixture; leaving magnesium in
solution. From solution countering alkali tartrate to complex the
aluminium, magnesium is precipitated as the quinolinate from
sodium hydroxide solution.
C) Analysis of mixture of aluminium and iron:
The iron is oxidised to the ferric state, then sodium hydroxide is
added in excess to precipitate iron as ferric hydroxide and
precipitate is filtered, washed, dissolved in nitric acid and
reprecipitated to obtain a purer precipitate of ferric hydroxide
and determined as under iron. The filtrate is acidified, and the
resulting aluminium salt is determined.
D) Analysis of mixture of iron and chromium:
The iron is first oxidised to the ferric state, and the chromium
is oxidised to the chromate state with sodium peroxide. The
containing the chromate, is acidified with acetic acid, and
treated with barium chloride. the precipitated barium
chromate is filtered, washed, dried, and weighed.
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