Chapter 3, Gravimetry

Gravimetric methods of analysis are used where weights of reactants and products of chemical reactions are reproducible, stable and reflect the presence of constituents which are important in the establishment of identity.

Two important methods deal with the trapping and weighing of products in the solid and gaseous phases. The first of these falls into the category of a precipitation method.

3-1. Precipitation methods. Many metallic elements in their ionic forms react with negative counter ions to produce stable precipitates. Silver ions form stable and highly insoluble salts with chloride, bromide and iodide. Calcium precipitates quantitatively with oxalate and can be measured reproducibly at any of three temperature dependent plateaus as the oxalate, the carbonate and the oxide. Barium precipitates quantitatively as the sulfate. The reactions often follow the same patterns:

Positive and negative ions in an aqueous solution, otherwise soluble with the counter ions in their environment, produce highly insoluble precipitates with certain added reagents.

3-2. Volatilization methods. An interesting volatilization method which is entirely gravimetric is the one shown by the equations below.

The analyte can be bicarbonate as shown or a mixture of carbonate and bicarbonate. The total amount of carbonate in whatever form is found by placing the analyte in a solution containing an excess of H2SO4. This solution is in a flask connected to incoming nitrogen gas gently bubbled through the solution and an exit tube first to a drying agent to absorb aerosolized water and water vapor and then to a mixture of NaOH and drying agent to absorb the CO2 and

water subsequently produced by the absorption by NaOH:

The apparatus is shown below. The tube containing the NaOH on asbestos and the CaSO4 to absorb the final water product is pre- and post-weighed to given the total amount of carbonate in the sample. Note that the nitrogen gas acts only as a carrier and does not take part in any reaction.

3-3. Considerations for the isolation of precipitates

Precipitates ought to be easy to wash free of contaminants without loss of the precipitate either in solution or through the filter. The particle size of the precipitate ought to be large enough not to escape through the filter pores. That the precipitate has a low solubility is paramount. The precipitate ought not to react with the atmosphere and it must have a known composition which remains stable after ignition.

3-3a. Substances of low solubility have the nasty habit of forming colloidal suspensions. Colloidal particles have diameters from 10-7 cm to 10-4 cm. That is on the order of from 10 atomic diameters to 10,000 atomic diameters. Particles in this size range are still sufficiently jostled about by thermal molecular motion to remain in suspension. Where they are the result of a process of precipitation brought about by the addition of ionic species, the particles are surrounded by the excess ionic species. If Ba2+ is added in excess to SO42- the BaSO4 precipitate which is formed is considered to be surrounded by Ba2+ ions. If the opposite procedure were being followed, the precipitate would be surrounded by SO42- ions. That these particles all have like charge and therefore repel each other suggests that your technique must favor the formation of large rather than small precipitate particles and to offer ways of encouraging the coagulation of particles after they have formed.

This can be done by carrying out the precipitation at a temperature close to the boiling point of water, in a dilute solution of your analyte and with constant stirring for the reasons given below.

Although analytical chemists still have some disagreement as to the mechanism of precipitation, there is wide agreement that a quantity called the relative supersaturation affects the particle size. Relative supersaturation is given as

where Q is the instantaneous concentration of the species added to effect precipitation and S is the equilibrium solubility of the substance which precipitates. Particle size seems to be inversely proportional to Relative Supersaturation because a high concentration of added reagent increases the probability that oppositely charged ions will begin the precipitation process at late as well as early stages of the addition and the resulting particles will be on the order of atomic dimensions, whereas the maintenance of a value of Q just slightly above S lowers that probability but offers in any case a layer of the added reagent ions around existing particles for their further growth.

3-4. The Electric Double Layer

If a particle of precipitate is surrounded by the ion in excess, say Ba2+ in the case of the determination of SO42-, any negative ions in the immediate surroundings will be attracted to that primary positive layer. In the case of the addition of a BaCl2 solution to a Na2SO4 solution, the ions Cl- and SO42- are available. As the sulfate is used up in the precipitation process it is the chloride which is left and which forms the second layer. Thus we have an electric double layer, made up first of barium ions then of chloride ions. This double layer keeps the colloidal precipitate particles from coming into contact with each other for further coagulation.

There are two ways to bring the particles closer together and to increase the probability of coagulation: (1) heating increases overall thermal motion, affecting both the mobility of adsorbed ions and of the colloidal precipitate particles themselves. The summary effect is that there are collisions of particles which result in the increase in particle size due to increased coagulation; (2) increasing the electrolyte concentration of the solution, for reasons not entirely clear, results in a decrease in the mean radius of the electric double layer and encourages further coagulation. Carrying out both of these operations results in digestion of the precipitate, an unfortunate term because biological digestion usually refers to the dissolving of food and absorption at the molecular level through the wall of the intestine. Digestion in quantitative analysis refers to the coagulation of a precipitate into a filterable form. Unfortunately after successful digestion, some of the primary electric layer is made up of Na+ ions which must be washed away ultimately for quantitative results to be achieved. The Ba2+ ions as well will give a positive error if not removed and end up being dried with the precipitate as excess BaCl2. Many coagulated precipitates do not respond well to washing with distilled water because as the second electric layer is removed (excess Cl- for example) the first remains on all particles with an electric charge of the same sign. The result is that there is a return to the repulsive state and an effective increase in the radius of the particles which then begin once again to separate as colloidal particles. The process is called peptization and is to be avoided if some of the precipitate is not to be lost. One way around this for many precipitates is to encourage digestion by heating and also by increasing the electrolyte concentration by washing with a reagent which will go off as a gas during the drying process. Dilute nitric acid, HNO3 , is effective for washing excess ions from AgCl. In choosing such a wash, it is imperative that the procedure has been carried out and has been shown to yield reproducible, quantitative results. Unexpected side reactions, complex formation and changes in solubility with added reagents are sufficiently unpredictable to make intuition in the absence of experience unacceptable.

3-5. Other demons which can plague quantitative precipitate isolation.

During the precipitation procedure a number of other problems can arise to give erroneous positive or negative results. Among these are surface adsorption, mixed crystal formation, occlusion and mechanical entrapment.

Any ions may be carried down during a precipitation as the result of surface adsorption. Na+ , or Cl- in the case of the determination of SO42- by the addition of dilute BaCl2 solution to a NaSO4 solution. Both Ba2+ and Na+ can compete for lattice positions as the particles form. Likewise, the ions Cl- and SO42- can have the same effect. In the quantitative determination of some transition metals, iron for example as Fe(OH)3, zinc, cadmium and manganese may be present as impurities and all three form sparingly soluble hydroxides as well, though each with greater solubility than the hydroxide of iron:

Compound Solubility Product
Fe(OH)3 4 x 10-38
Cd(OH)2 2.5 x 10-14
Mn(OH)2 1.9 x 10-9
Zn(OH)2 1.2 x 10-17

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

Interfering Ions
K+, 133 pm NH4+, 148 pm
Sr2+, 113 pm Ba2+, 135 pm
Mn2+, 80 pm Cd2+, 97 pm

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

Occlusion and mechanical entrapment. If a precipitation procedure is carried out too quickly, pockets of solvent and spectator ions can form, trapping them within the precipitate particles and dashing one's hope of removing them during the washing procedure. This is another reason why the relative supersaturation must be kept as low as possible so that in principle at least, all precipitation occurs only at the surface of a growing solid particle, devoid of solvent pockets.

All of these problems of coprecipitation of unwanted ions can lead to positive or negative errors. In the example above where it pointed out that Na+ or Cl- may coprecipitate in the SO42- determination, surface adsorption will produce a positive error. In the case of mixed crystal formation, the direction of the error depends on the relative atomic weight of the ion which replaces that which is desired in the precipitate. In the case of the precipitation of zinc hydroxide, mixed crystal formation with manganese would produce a negative error but with cadmium or zinc a positive error.
Desired precipitate
Compound Mn(OH)2 Zn(OH)2 Cd(OH)2
At. Wt. of M2+ 54.94 65.39 112.41
Direction of error negative --- positive

3-6. The use of the technique of homogeneous solutions to effect precipitation.

A solution containing a reagent which produces a desired ion to effect precipitation, often by gentle heating of the solution, offers an exquisite means for obtaining well-formed large crystal particles which lend themselves splendidly to the technique of filtration.

The model we use to explain why this happens also uses the concept of relative supersaturation. The initial nucleation of sparingly soluble particles offers a surface template which favors "locking" onto ions in the vicinity which by the luck of the draw (and the kinetic molecular theory) find themselves at the right energy and orientation to enter the crystal lattice. Ions isolated from a growing crystal are not favored to enter this process because at least two are required, both at the right energy and orientation to start the growth of a new crystal. If the concentration of one ion of a sparingly soluble salt increases gradually by slow homogeneous synthesis in a solution, then as its concentration reaches the threshold of supersaturation for the ion pair, a relatively small number of nucleated particles grows to larger size (because the probability of finding a place in an existing crystal lattice for any single ion is greater than that of a spontaneous creation a new crystal from dissolved and randomly arranged ions) rather than a large number of nucleated particles growing in constant competition with the rest and thus remaining small. The result for the latter is a non-filterable precipitate, but one in the former which filters quite well.

See the demonstration of this effect at

Here is a table of common reagents useful for the preparation of ions often needed for precipitation processes.

Reagent Precipitating species Precipitation reaction Elements which yield to

this reaction

Urea OH- (NH2)2CO + 3H2O --->

CO2 + 2NH4+ + 2OH-

Al, Ga, Th, Bi, Fe, Sn
Trimethyl phosphate PO43- (CH3O)3PO + 3H2O ---> 3CH3OH + H3PO4 Zr, Hf
Ethyl oxalate C2O42- (C2H5)2C2O4 + 2H2O ---> 2C2H5OH + H2C2O4 Mg, Zn, Ca
Dimethyl sulfate SO42- (CH3O)2SO2 + 4H2O ---> 2CH3OH +SO42- + 2H3O+ Ba, Ca, Sr, Pb
Trichloroacetic acid CO32- Cl3CCOOH + 2OH- ---> CHCl3 + CO32- + H2O La, Ba, Ra
Thioacetamide H2S CH3CSNH2 + H2O --->


Sb, Mo, Cu, Cd
Dimethyl glyoxime CH3(CNOH)2CH3 CH3COCOCH3 +

2H2NOH ---> DMG + 2H2O

8-Acetoxyquinoline C9H6NOH CH3COOQ + H2O ---> CH3COOH + HOQ Al, U, Mg, Zn

3-7. Preparation of a dry weight of your precipitate.

The resulting precipitate must be heated until a stable dry state is reached. Some understanding of typical precipitate properties is mandatory for repeatable results to be achieved.

Note in the figure at the right that whereas AgCl achieves a stable dry weight just above 100oC, BaSO4 does not do so until it reaches a temperature in the vicinity of 700oC Aluminum oxide, Al2O3, loses water slowly as the temperature rises to 1000oC at which point it achieves stability. Some compounds decompose in several stages, reaching stable plateaus. Calcium oxalate, CaC2O4 H2O, loses all its water at around 200oC and remains stable as CaC2O4 until just above 400oC at which point it decomposes to calcium carbonate, CaCO3 where it remains stable up to 700oC. Between 700oC and 850oC it slowly decomposes to CaO where it remains stable until its melting point at 2614oC.

A device not seen often in analytical laboratories but useful for producing automatic plots of mass of sample vs. temperature such as those at the right is the thermobalance (below). Region A includes the heating circuit, a temperature sensor, sample cup and counter weight resting on one end of the balance arm, B. A light wave is partially attenuated at C, giving a negative feedback to the amplifier circuit at D, designed to yield an output voltage which increases with the force necessary to keep the balance in equilibrium (and the attenuation at a constant value). One can adjust the baseline voltage at E, the tare adjuster, so as to produce the graph at the chart recorder, F.

Example 3-1: A 0.3427 g sample of bronze-age jewelry is analyzed for silver content by first dissolving it in concentrated nitric acid and precipitating it as AgCl. The precipitate is transferred to a dry sintered glass filter weighing 12.2347 g where it is separated from the filtrate and washed with dilute nitric acid. The filter and precipitate are dried at 150oC, cooled, weighed and found to weigh 12.4373 g. Calculate the %Ag in the jewelry.

(To be solved and discussed in class)

Example 3-2. A sample of iron ore weighing 0.4275 g is dissolved in 12M HCl. The resulting solution is slowly made basic with NaOH until the first hint of a turbid solution is detected. 10.0 g urea are dissolved and the solution heated just to the boiling point for 4 hours. The precipitate, Fe2O3 x H2O, is trapped using ashless filter paper. The precipitate and filter paper are fired in a porcelain crucible of empty weight 12.2837g until the ashless filter paper is completely incinerated and anhydrous Fe2O3 is left. The resulting weight of crucible and precipitate is 12.4274 g. Determine the %Fe, the %Fe2O3 and the %Fe3O4.

(To be solved and discussed in class)

Example 3-3. A sample known to contain only KCl and NaCl and weighing 0.4263 g is dissolved in water and treated to an excess of AgNO3, using standard methods of precipitation. The AgCl precipitate is caught on a Gooch Crucible of original dry weight of 15.2748 g. The AgCl precipitate is dried at 150oC, cooled and the crucible and precipitate are found to weigh 16.2872 g. Determine the %KCl and the %NaCl in this sample.

Solution is based on the difference in the %Cl in the pure salts:

Salt %Cl
NaCl 60.66
KCl 47.55

(To be solved and discussed in class)

3-8. Preferred methods of gravimetric analysis. Most inorganic ions have yielded to gravimetric analytical techniques, but one finds many interfering ions. The table below illustrates both the abundance of reagents available for use as well as the problems which can be encountered by interfering ions:

Analyte Precipitate Measured form Interferences
K+ KB(C6H5)4 KB(C6H5)4 NH4+,Ag+,Hg2+, Tl+,Rb+,Cs+
Mg2+ Mg(NH4)PO4.6H2O Mg2P2O7 Many metals (none from Na+ and K+)
Ca2+ CaC2O4.H2O CaCO3 or CaO Many metals (none from Mg2+, Na+ and K+)
Ba2+ BaSO4 BaSO4 Na+,K+,Li+,Ca2+,Al3 + ,Cr3+,Fe3+,Sr2+,Pb2+
Ti4+ TiO(5,7-dibromo-8-hydroxyquinoline)2 TiO(5,7-dibromo-8-hydroxyquinoline)2 Fe3+,Zr4+,Cu2+,C2O4 2-, citrate, HF
VO43- Hg3VO4 V2O5 Cl-,Br-,I-,SO42- , CrO42-,AsO43-,PO43-< / sup>
Cr3+ PbCrO4 PbCrO4 NH4+,Ag+
Mn2+ Mn(NH4)PO4.H2O Mn2P2O7 Interferences from numerous metals
Fe3+ Fe(HCO2)3 Fe2O3 Interferences from numerous metals
Co2+ Co(1-nitroso-2-naphtholate)3 CoSO4 (by reaction with H2SO4 ) Fe3+,Zr4+,Pd2+
Ni2+ Ni(dimethylglyoximate)2 Ni(dimethylglyoximate)2 Pd2+,Pt2+,Bi3+,Au3+
Cu2+ CuSCN CuSCN NH4+,Pb2+,Hg2+,Ag+
Zn2+ Zn(NH4)PO4.H2O Zn2P2O7 Interferences from numerous metals
Ce4+ Ce(IO3)4 CeO2 Th4+,Ti4+,Zr4+
Al3+ Al(8-hydroxyquinolate)3 Al(8-hydroxyquinolate)3 Interferences from numerous metals
Sn4+ Sn(cupferron)4 SnO2 Cu2+,Pb2+,As(III)
Pb2+ PbSO4 PbSO4 Ca2+,Sr2+,Ba2+,Hg2+, Ag+,HCl, HNO3
NH4+ NH4B(C6H5)4 NH4B(C6H5)4 K+, Rb+, Cs+
Cl- AgCl AgCl Br-, I-, SCN-, S2-, S2O32-, CN-
Br- AgBr AgBr Cl-, I-, SCN-, S2-, S2O32-, CN-
I- AgI AgI Br-, Cl-, SCN-, S2-, S2O32-, CN-
SCN- CuSCN CuSCN NH4+,Pb2+,Hg2+,Ag+
CN- AgCN AgCN Cl-, Br-, I- , SCN-, S2-, S2O32-
F- (C6H5)3SnF (C6H5)3SnF Except alkali metals, many interferences,

and SiO44- , CO32-

ClO4- KClO4 KClO4
SO42- BaSO4 BaSO4 Na+,K+,Li+,Ca2+,Al3 + ,Cr3+,Fe3+,Sr2+,Pb2+
PO43- Mg(NH4)PO4.6H2O Mg2P2O7 Many interferences except Na+,K+
NO3- Nitron nitrate Nitron nitrate ClO4-, I-, SCN-, CrO42-,ClO3-, NO2-, Br-, C2O42-
CO32- CO2 (by addition of acid) CO2 CO2 is trapped as Na2CO3 on Ascarite

There are a number of organic functional groups which precipitate with metal ions by one of two routes: (1) chelating agents are organic compounds which "wrap around" a metal ion thanks to cationic side chains which form coordinate covalent bonds with the ion, and (2) a straightforward ion-ion bond which produces a new species that excludes water of solvation and thus precipitates. Good examples of chelating agents include Ethylene Diamine Tetraacetic Acid (EDTA), oxalic acid, glycine, 8-hydroxyquinoline and dimethylglyoxime.

Some common organic precipitating agents:

Compound Ions precipitated
Dimethylglyoxime Ni2+,Pd2+,Pt2+
EDTA (Ethylenediamine tetraacetic acid) Zn2+, Cu2+, Pb2+, Ca2+, Ni2+, Fe3+
Cupferron Fe3+,VO2+,Ti4+, Zr4+,Ce4+,Ga3+,Sn4+
8-Hydroxyquinoline Fe3+,Al3+,Mg2+,Zn2+, Cu2+,Cd2+,Pb2+, Bi3+, Ga3+,Th4+, Zr4+, TiO2+, UO22+
Salicylaldoxime Bi3+,Ni2+,Pd2+,Zn2+, Cu2+,Pb2+
1-Nitroso-2-naphthol Fe3+,Co2+,Pd2+, Zr4+
Nitron (C20H16N4) NO3-, ClO4-, BF4-, WO42-
Sodium tetraphenylborate NH4+, organic ammonium, Ag+, Cs+, Rb+, K+
Tetraphenylarsonium chloride Cr2O72-, MnO4-, ReO4-, MoO42-,WO42-, ClO4-