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3. Gravimetry
Gravimetry includes any analytical method in which the ultimate measurement is
by weight.
Steps of gravimetric analysis:
The individual steps required in a gravimetric analysis can be summarized as follows:
1)
2)
3)
4)
Preparation of the solution
precipitation
digestion
filtration
5)
6)
7)
8)
washing
drying or igniting,
weighing,
calculation.
1) Preparation of the solution:



Eliminate interfering materials.
Proper adjustment of the solution conditions prior to precipitation
Factors that must be considered include the volume of the solution during
precipitation, the concentration range of the test substance, the presence
and concentrations of other constituents, the temperature, and the pH.
The pH is important because incase of the possibility of interferences from other
substances. For example:
 calcium oxalate is insoluble in basic medium, while in acid medium
the
oxalate ion combines with the hydrogen ions to form a weak acid,

8-Hydroxyquinoline (oxine) can be used to precipitate a large number of
elements, but by control of pH, elements can be precipitated selectively.
Conditions for analytical precipitation:
1.
The precipitate should consist of large crystals so that they can be easily
filtered.
2. This contamination should be negligible; it is minimized also by keeping the
crystals large.
The precipitation occurs in a series of steps.
1. First, supersaturation occurs, that is, the solution contains more of the dissolved
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salt than occurs at equilibrium.. This is started by nucleation, number of particles
must come together to produce microscopic nuclei of the solid phase. The higher
the degree of supersaturation, the greater the rate of nucleation.
Following nucleation, the initial nucleus will grow by deposition of other precipitate
particles to form a crystal of a certain geometric shape. Again, the greater the
supersaturation, the more rapid the crystal growth rate. With increased growth rate,
there is increased chance of imperfections in the crystal and trapping of impurities.
Von Wiemarn discovered that the particle size of precipitates is inversely
proportional to the relative supersaturation of the solution during the precipitation
process:
relative supersaturatiou =
QS
S
where Q is the concentration of the mixed reagents before precipitation occurs and is
the degree of supersaturation, and S is the solubility of the precipitate at equilibrium.
This ratio is also called the von Wiemarn ratio.
high relative supersaturation → many small crystals (high surface area)
low relative supersaturation → fewer large crystals (low surface area)
Obviously, then, we want to keep Q low and S high during precipitation. Several steps
are commonly taken to maintain favorable conditions for precipitation.
1- Precipitate from dilute solution. This keeps Q low.
2- Add dilute precipitating reagents slowly, with effective stirring. This also keeps
Q low.
3- Precipitate from hot solution. This increases S.
4- Precipitate at as low a pH as is possible to still maintain quantitative
precipitation.
Most of these operations also often decrease the degree of contamination. The
concentration of impurities is kept lower and their solubility is increased, and the
slower rate of precipitation decreases their chance of being trapped. The large
crystals have a smaller surface area relative to the mass and so have less chance of
adsorption of impurities.
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2) Digestion of the precipitate:
When a precipitate is allowed to stand in the presence of the mother liquor (the
solution from which it was precipitated), the large crystals grow at the expense of
the small ones. This process is called digestion, or Ostwald ripening. The small
particles tend to dissolve and precipitate on the surfaces of the larger crystals. In
addition, individual particles agglomerate to effectively share a common counter-ion
layer and the agglomerated particles finally cement together through formation of
connecting bridges. This result is an appreciable decrease in surface area.
Also, imperfections of the crystals tend to disappear, and adsorbed or trapped
impurities tend to go into solution. It improves the filterability of the precipitate and
its purity.
Many precipitates do not give a favorable von Wiemara ratio, especially very
insoluble ones. Hence, it is impossible to yield a crystalline precipitate (small number
of. large particles) and the precipitate is first colloidal (large number of small
particles).
Colloidal particles are very small (1-100 m) and have a very large surface which
promotes surface adsorption. In AgCl, for example, there will be alternating Ag+ and
Cl- ions on the surface. While there are localized + and - charges on the surface, the
net surface charge is zero. The adsorption creates a primary layer that is strongly
adsorbed and is an integral part of the crystal. It will attract ions of the opposite
charge in a counterlayer or secondary layer and give an overall neutral particle.
The primary surface charge will
tend to repel like particles,
creating a colloidal state. An
added electrolyte will result in
a closer secondary layer and
will promote coagulation.
Heating tends to decrease
adsorption and the effective
charge in the adsorbed
layers,
thereby
aiding
coagulation. Stirring will also
help.
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There are two types of colloids, hydrophilic and hydrophobia Hydrophilic means
"water-loving" and these colloids have a strong affinity for water. A solution of a
hydrophilic colloid is therefore viscous. A hydrophobic colloid has little attraction
for water. A solution of this type of colloid is called a sol. Coagulation of a hydrophobic
colloid is fairly easy and results in a crudy precipitate.
An example is silver chloride. Coagulation of a hydrophilic colloid, such as hydrous
ferric oxide, is more difficult, and it produces a gelatinous precipitate that is
difficult to filter because it tends to clog the pores of the filter. In addition,
gelatinous precipitates adsorb impurities readily, because of their very large
surface area. Sometimes a reprecipitation of the filtered precipitate is required. To
reduce concentration of impurities.
Impurities in precipitates: (contamination)
Precipitates tend to carry down from the solution other constituents that are
normally soluble, causing contamination of the precipitate. This process is called either,
coprecipitation or postprecipitation.
There are a number of ways in which a foreign material may be coprecipitated.
1. Occlusion. foreign ions are trapped inside the
crystal as it is formed. For example, when BaSO4 is
precipitated by the slow addition of Ba2+ (as BaCl2) to a
solution of SO42-, the sulfate ion will be in excess
throughout the precipitation, and the primary
adsorbed layer will be sulfate ions. This is neutralized
by a secondary layers of counter ions from the solution. For example, Na+ may be the
counter ion (two per SO42-). As more barium chloride is added, the loosely bound Na+
is replaced by Ba2+, and the crystal continues to grow. However, if the growth takes
place too rapidly, all the sodium ions may not be replaced by barium ions.
Occluded impurities are difficult to remove. Digestion may help some, but this is not
completely effective. The impurities cannot be removed by washing. Purification by
dissolving and reprecipitating may be helpful.
2. Surface adsorption. the surface of the precipitate will have a primary adsorbed
layer of the lattice ions in excess. After the barium sulfate is completely
precipitated, the lattice ion in excess will be barium, and this will form the primary
layer. The counter ion will be foreign anion, say, a nitrate anion, two for each barium
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ion. The net effect is an adsorbed layer of barium nitrate. These adsorbed layers can
often removed by washing, or they can be replaced by ions that are readily
volatilized. Gelatinous precipitates are especially troublesome, though digestion
reduces the surface area and the amount of adsorption.
3. Isomorphous replacement. Two compounds are said to be isomorphous if they
have:
1. The same type of formula
2. Crystallize in similar geometric forms.
3. Same lattice dimensions.
one ion can replace another in a crystal, resulting in a mixed crystal. For example, in
the precipitation of Mg2+ as magnesium ammonium phosphate, K+ has nearly the same
ionic size as NH4+ and can replace it to form magnesium potassium phosphate.
Isomorphous replacement is very serious and little can be done about it.
4. Postprecipitation. Sometimes, when the precipitate is allowed to stand in contact
with the mother liquor, a second substance will slowly form a precipitate with the
precipitating reagent. For example, when calcium .oxalate is precipitated in the
presence of magnesium ions, magnesium oxalate does not immediately precipitate.
But it will come down if the solution is allowed to stand too long before being filtered.
Washing and filtering the precipitates:
Coprecipitated impurities, especially those on the surface, can be removed by
washing the precipitate after filtering. The precipitate will be wet with the mother
liquor, which is also removed by washing. Many precipitates cannot be washed with
pure water, because peptization occurs. This is the reverse of coagulation.
The process of coagulation discussed above is
reversible. As mentioned, coagulated particles
have a neutral layer of adsorbed primary and
counter ions. Washing with water will dilute the
counter layer, developing repulsive forces
between individual particles. The coagulated
particles thereby peptize and form colloidal
particles that pass through the filter. Prevention consists in adding an electrolyte
to the wash liquid.
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The electrolyte must be one that is volatile at the temperature to be used for drying or
ignition, and it must not dissolve the precipitate. For example, dilute nitric acid is
used as the wash solution for silver chloride. The nitric acid replaces the adsorbed
layer of Ag+ /anion-, and it is volatilized when dried at 110°C. Ammonium nitrate is used
as the wash electrolyte for hydrous ferric oxide. It is decomposed to NH3, HNO3,
N2, and oxides of nitrogen when the precipitate is dried by ignition at a high
temperature.
When a precipitate is washed, a test should be made to determine when the washing
is complete. This is usually done by testing the filtrate for the presence of an ion
of the precipitating reagent. After several washings with small volumes of the wash
liquid, a few drops of the filtrate are collected in a test tube for testing. For
example, if chloride ion is determined by precipitating with silver nitrate reagent,
the filtrate is tested for silver ion by adding sodium chloride.
Drying or ignition of the precipitate:
The precipitates may contain:
1. adsorbed water(water of humidity)
2. Occluded water (present in cavities within crystals).
3. sorbet water, with substances having a large internal surface
4. essential water, as water of crystallisation [e.g. CaC204.H2O
Simply drying can usually be done by heating at 110°-120°C for one to two hours.
Ignition at a much higher temperature is usually required For example, magnesium
ammonium phosphate (MgNH4PO4) is decomposed to the pyrophosphate (Mg2P2O7) by
heating at 900°C.
In addition to the evolution of water, the ignition of precipitates often results in
thermal decomposition reactions involving the dissociation of salts into acidic and
basic components, e.g. decomposition of carbonates and sulphates. The temperatures
at which precipitates may be dried, or ignited to the required chemical form, can be
determined from a study of the thermogravimetric curves for the individual
substances.
Thermogravimetry is a technique in which a change in the weight of a substance is
recorded as a function of temperature or time. The results presented as a
thermogravimetric (TG) curve, in which the weight change is recorded as a function
of temperature or time,
A typical thermogravimetric curve, for copper sulphate pentahydrate CuSO3.5H2O, is
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given in the next Figure.
Homogeneous precipitation:
We have seen that the most favorable conditions for precipitation involve
precipitating from dilute solution and adding the precipitating agent slowly with
effective stirring. These operations maintain a low degree of supersaturation (low
Q). local excesses of precipitating reagent are unavoidable when the reagent is
added to the sample solution. A process called homogeneous precipitation avoids
such difficulties.
In this technique, the precipitating reagent is generated by a chemical reaction which
occurs uniformly throughout the solution.
An example is the generation of hydroxyl ions for the precipitation of hydrous iron
oxide, aluminum oxide, and so forth. The precipitating reagent is formed by hydrolysis
of urea in an initially acid solution:
(NH2)2CO + 3 H2O → CO2  + 2 NH4+ + 2OHThis hydrolysis reaction takes place slowly and occurs at temperatures just below the
boiling point of water.
The precipitate formed by this technique:
1. Have higher densities and much more easily filtered.
2. The degree of coprecipitation is decreased and a purer precipitate is obtained.
Examples of homogeneous precipitation reactions, Sulfate ion can be generated
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homogeneously by heating a solution containing sulfamic acid. The sulfamic acid
hydrolyzes as follows:
HSO3NH2 + H2O → H+ + SO42Thus, barium sulfate or lead sulfate may be homogeneously precipitated. In the
presence of nitric acid, sulfate ion is also formed homogeneously by the following
reaction:
HSO3NH2 + HNO3 →2H+ + SO42- + N2O + H2O
Calculations:
Calculations are usually made on a percentage basis:
A% 
gA
grams of analyte
x100 
x100%
AS
grams of sample taken for analysis
For example, if we are analyzing for the percentage of chloride in a sample by
weighing it as AgCl, we can write:
precipitating agent
Cl  AgCl
We derive one mole of AgCl from one mole of Cl-, so
grams Cla.w.Cl
a.w.Cl
and g Cl  g AgCl x

F.w.AgCl
grams of AgCl F.w.AgCl
If we precipitate the chloride as PbCl2,
precipitating agent
2 Cl        PbCl
and
g Cl  g pbCl2 x
2 a.w.Cl
F.w.pbCl2
then
grams Cl2 a.w.Cl

grams of PbCl2 F.w.pbCl2
In general, the gravimetric factor is the appropriate ratio of the formula weight of
the substance sought to that of the substance weighed:
Gravimetri c factor =
f.w (substance sought)
f.w. (substance weight)
In the above example, the gravimetric factors were Cl/AgCl. and 2Cl/PbCl2.
The weight of the substance sought is obtained by multiplying the weight of
the precipitate by the gravimetric factor:
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We can now write a general formula for calculating the percentage of
composition of the substance sought:
% sought=
weight of precipitate (g) x gravimetric factor
weight of sample (g)
X 100
Example. Orthophosphate (PO43-) is determined by weighting as ammonium
phosphomolybdate, (NH4)3PO4.12MoO3. Calculate the percent of P in. the sample and
the percent of P2O5 if the precipitate from a 0.2711-g sample weighed 1.1682g.
Solution:
1.168 x [p/(NH4) 3PO4 .12MoO 3]
x100%
0.2711
1.168 x30.97/187 6.5
%p
x100%  7.111%
0.2711
%p
1.168 x [p2O5/(NH4)3 PO4 .12MoO 3 ]
x100%
0.2711
1.168 x [141.95/2x 1876.5 ]
% p2O5 
x100%  16.30%
0.2711
% p2O5 
Organic precipitates:
Organic precipitating agents have the advantages of giving precipitates with:
1. very low solubility in water
2. Favorable gravimetric factor. (more sensitive and precise)
Most of them are chelating agents that form slightly soluble uncharged chelates
with the metal ions.
Mn+ + nHX → MXn + nH+
For increasing the selectivity the pH must adjusted
Dimethylglyoxime with Nickel, lead , platinum
cupferron with Iron, stannous ….
8-hydroxyquinoline (oxine)
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1. chloride as AgCl
2. sulphate as BaSO4:
3. iron as iron(III) oxide
Precipitation
by precipitation of Cl- as its
silver salt after acidification with
nitric acid.
Ag+ + Cl- →AgCl
based
upon
based upon the precipitation of ferric iron as a
precipitation with Ba++ ion:
hydrous oxide,
Ba++ + SO42- →BaSO4 2Fe+++ + 6NH3 + (x+3)H2O → Fe2O3.xH2O + 6NH4+
ammonium hydroxide is used as precipitant. The
precipitate (amorphous) forms colloidal solution. To avoid
this, the precipitation must be done in hot solution in
presence of a coagulating electrolyte.
Washing
The precipitate washed with very
dilute nitric acid, in order to
prevent peptisation,
washed with water
The most satisfactory wash liquid is ammonium nitrate;
this voltatilizes upon ignition. Ammonium chloride is
unsuitable, because iron(III) chloride which is volatile, is
formed during the ignition.
Fe2O3 + 6NH4Cl = 2FeCl3 + 6 NH3 + H2O
Drying and
igniation
Dried at 130-150°C, and then
finally weighed as AgCl.
ignited at temperature less
than 600oC and weighed as
BaSO4
by ignition Fe2O3.xH2O to Fe2O3.
Fe2O3.xH2O → Fe2O3 + xH2O
precautions
i) Precipitation is carried out in  BaSO4 ia a very small crystals, which sometimes
nitric acid solution in order to
pass through the filter pores pores. To overcome
prevent the precipitation of
this problem:
other silver salts, such as a) The precipitant should be added slowly.
phosphates and carbonates.
b) Addition of substance that increases the
ii) excess of Ag+ ion is required to
solubility of the precipitate (decreased the
decrease the solubility of the
degree of supersaturation ). Hydrochloric acid can
precipitate.
play this role by the formation of HSO42-.
iii) The precipitate of AgCl must
c) The precipitation at boiling temperature.
be protected from light; to
prevent photochemical
 Pure BaSO4 is not decomposed when heated in dry
decomposition of the precipitate
air until a temperature of about 1400°C is reached:
occurs.
BaSO4 →BaO + SO3
39
i) Cations as Al+++, Cr+++, Mn3+ ,
anions arsenate, phosphate, etc.
must be also absent.
ii) Organic hydroxy compounds
e.g. citric acid, tartaric acid,
sugar,
etc.
prevent
the
precipitation of iron by forming
complex salts.
iii) Cupferron can be used as a
specific one for precipitation of
ion. The precipitated complex
has also to be ignited to ferric
oxide (the common weighying
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AgO (Violet color) + ½ Cl2
(
iv) I-, Br-, and SCN-, interfere.
AgCl

The precipitate is, however, easily reduced to
sulphide at the temperature above 600°C by the
carbon of the filter paper.
BaSO4 + 4C → BaS + 4 CO
form of iron). Iron can also be
precipitated as oxinate which is
filtered, washed, dried and
weighed.
4. aluminium as aluminium oxide:
5. magnesium:
Precipitation
The same way as Fe+++, by precipitation
with ammonia, and subsequent
conversion of the A1(OH)3 precipitate
into Al2O3 by ignition.
diammonium hydrogenphosphate, and excess of ammonia solution is added to
precipitate ammonium magnesium phosphate hexahydrate, Mg(NH4)PO4.6H2O at room
temperature.
Mg2++ HPO42- + NH4+ + OH- = Mg(NH4)PO4 + H2O
The precipitate should stand for several hours before filtration (to decrease
its solubility).
Washing
The precipitate is washed with either
ammonium chloride or ammonium
nitrate solution; to prevent
peptisation.
The precipitate is washed with 0.8M aqueous ammonia solution and then
weighed either as the hexahydrate MgNH4PO4.6H2O or as the pyrophosphate
Mg2P2O7.
For the former, the precipitate is washed with ethanol, followed by anhydrous
diethyl-ether, and weighed after standing in a desiccators for about 20 minutes at
room temperature. This method is of moderate accuracy and is recommended
only because of the time saving achieved.
Drying and
igniation
Al2O3, obtained after ignition is hygroscopic unless the
temperature of ignition is raised to at least 1200°C at which a non
hygroscopic form of Al2O3 is formed. A13+ may also be precipitated
with 8-hydroquinoline at pH 4-10.
The precipitate is ignited at a high temperature
and weighed as such.
2MgNH4PO4.6H20 = Mg2P2O7 + 2NH3 + 13H2O
A13+ + 3H(C9H6NO) → A1(C9H6NO)3 + 3H+
For complete precipitation of Al+++ the pH of the solution must be carefully controlled. To about 6.5-7.5. The addition of
ammonium chloride exerts a buffering effect and also assists the coagulation of the colloidal precipitate.
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6. Mixture containing 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 convert aluminium into soluble aluminate A1O2-.
The precipitate is filtered, washed, dissolved in nitric acid and reprecipitated to obtain a purer
precipitate of ferric hydroxide and determined as before. The filtrate is acidified, and the resulting
aluminium salt is determined as explained before.
7 Determination of Calcium:
The calcium is precipitated as CaC2O4.H2O by treating hot acidified solution with ammonium oxalate, and slowly
neutralizing with aqueous ammonia solution.
Ca2+ + C2O42- + H2O → CaC2O4.H2O
The precipitate is washed with dilute ammonium oxalate solution and then weighed in one of the following
forms:
1. As CaC2O4.H2O by drying This method is not recommended because of the hygroscopic nature of
at 100-105°C for 1-2 the oxalate and the difficulty of removing the co-precipitated ammonium
hours.
oxalate at this low temperature.
2. As CaCO3 by heating at
about 500°C
since the CaCO3 is not hygroscopic, this is considered as the most
satisfactory method:
Ca2C2O4 → CaCO3 + CO
3. As CaO by ignition at
1200°C.
This method is widely used, but the resulting CaO has a
comparatively small molecular weight and is hygroscopic.
CaCO3 → CaO + CO2
Pyrolysis curve (relation between mass and temperature)
To obtain course crystal ammonium oxalate , followed by ammonium hydroxide to a hot acid solution
of the calcium salt drop by drop


Removal of the Cl- ions by washing prevents losses due to the formation of volatile CaCl2
during ignition of the precipitate.
All metals which form slightly soluble oxalate must be absent e.g. Cu2+, Pb2+ and Zn2+.
41